George Samuel Newth.

A manual of chemical analysis, qualitative and quantitative online

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compounds in anything approaching to a condition of purity, while with most
of the metals the cost of the salts practically prohibits such a course (see note,
p. 26). These elements, however, are not all equally " rare," and therefore in
this section some of the characteristic reactions of a few of the most commonly
occurring of these metals are given.

f Thallium also appears in Group I. (see p. 16), since the chloride is
precipitated by hydrochloric acid from thalloits solutions; and, like lead
chloride, is slightly soluble in water.

% See footnote on p. 18.

The Rare Metals of Group III 71

regarded as belonging to the same natural family. Thus, ammonia,
potash, and soda produce, with solutions of either metal, a white
flocculent precipitate of the respective hydroxides, A1,(HO) C and
Be(HO) 2 .

Both precipitates are soluble in excess of the fixed alkalies, but
the beryllium compound is reprecipitated if the diluted solution is

Beryllium hydroxide is decomposed by boiling with ammonium
chloride, ammonia being evolved and beryllium chloride passing
into solution ; thus

Be(HO) 2 + 2NH 4 C1 = BeCl 2 + 2NH 3 + 2H 2 O

Aluminium hydroxide undergoes no change when similarly treated,
hence the metals may be separated by this reaction.

The carbonates of the alkalies give precipitates of beryllium
carbonate, or basic carbonate ; readily soluble in excess of ammo-
nium carbonate, giving a double carbonate.

The precipitate is soluble, but far less readily, in the carbonates
of the fixed alkalies. Beryllium is therefore readily separated from
aluminium by adding ammonium carbonate and warming the
solution. Aluminium hydroxide is precipitated, while the beryllium
goes into solution as the double carbonate of ammonium and
beryllium. On filtering off the aluminium hydroxide, the beryllium
hydroxide may be precipitated by ammonia, after first neutralising
with hydrochloric acid.

The salts of beryllium possess a characteristic sweet taste,
hence the name ghtcinum, which was formerly applied to the

Zirconium. The oxide of this element, along with others of
the so-called " rare earths " (but more especially Zirconia), has the
property of remaining unchanged for a long time when heated to
incandescence, and of emitting a bright white light when so heated.
It is on this account used, with others of the rare earths, in the
construction of the " mantles " of the incandescent gas-burners
now so common.

Alkaline hydroxides, as well as ammonium sulphide (group-
reagent), give a white precipitate of zirconium hydroxide, Zr(HO) 4 ,
resembling aluminium hydroxide in appearance. It is distinguished
from the latter, in that it is insoluble in excess of potassium or
sodium hydroxide.

Alkaline carbonates give a white precipitate consisting of a
basic carbonate, which is soluble in excess, especially of ammonium

Potassium sulphate (a concentrated boiling solution) gives a
precipitate of a double sulphate, which, when thrown down from
the hot solution, is scarcely soluble in hydrochloric acid (thorium
gives a similar precipitate under the same conditions, but the
thorium compound is soluble in hydrochloric acid). In dilute
solutions the precipitate only appears after standing for some hours.
Sodium thiosulphate gives a precipitate of zirconium thiosulphate,

72 Qualitative Analysis.

which on boiling is complete even in very dilute solutions (thorium
behaves similarly). Oxalic acid gives a white precipitate of
zirconium oxalate, soluble in ammonium oxalate. (Thorium oxalate,
similarly precipitated, is insoluble in ammonium oxalate.)

Hydrogen peroxide, added to a slightly acid solution of a
zirconium salt, gives a white precipitate, believed to be either ZrO 3
or Zr 2 O 5 . (Niobium and titanium do not give a precipitate.)

By means of these three reactions zirconium can be separated
from the other " rare earths " of this group.

The hydrated oxides, which are precipitated by ammonia, are
first dissolved in hydrochloric acid, and the solution nearly
neutralised with sodium carbonate. Sodium thiosulphate is then
added, and the mixture boiled. The precipitate may consist of the
thiosulphates of zirconium and thorium, together with titanic acid.
The precipitate is treated with boiling hydrochloric acid, which
dissolves the two thiosulphates, and possibly a little titanic acid.
Excess of ammonium oxalate is added to the solution, which at
first precipitates oxalates of zirconium and thorium, but redissolves
the zirconium oxalate. The insoluble thorium oxalate is removed
by filtration, and any titanium in the solution is precipitated by
means of ammonium carbonate, added in excess in order to re-
dissolve the zirconium basic salt, which is first thrown down.
Any permanent precipitate is filtered off, and the filtrate con-
centrated by gentle evaporation. A boiling strong solution of
potassium sulphate is then added, which precipitates the double
sulphate of potassium and zirconium.*

Titanium. This element is met with in small quantities in
many specimens of iron ores, clays, and igneous rocks. The
minerals special to the element, rutile^ anatase (TiO 2 ) ; sphetie
(silicate and titanate of calcium), are rare substances titaniferous
iron (ferrous titanate) is less rare.

Titanium oxide, TiO 2 , being insoluble in hydrochloric and
nitric acid, is found in the insoluble residue after treatment with
acids ; titanates, on the other hand, are dissolved by hydrochloric
acid, but on boiling the solution white titanic acid, H 2 TiO. 3 , is

Titanium oxide (in the absence of other metals which colour a
bead of microcosmic salt), when heated in a bead of microcosmic
salt in the inner blowpipe flame, imparts to the bead a colour
which is yellowish when hot, but which becomes violet as the bead

* This separation, and the reactions for zirconium, may be made by dis-
solving up a couple of the incandescent gas "mantles ; " either new ones, or
those which have become worn out by use. In the former case they should
first be set fire to, in order to burn off the organic matter present. The
mantles are boiled in a test-tube with a little strong sulphuric acid, and the
liquid when cold diluted with water. The insoluble residue (consisting of
the main portion of the material) is filtered off, and the hydrated oxides of the
rare earths, along with any alumina, are precipitated by the addition of
ammonia. The precipitate is then washed, and dissolved in a small quantity
of hydrochloric acid.

The Rare Metals of Group III. 73

cools. The colour is more readily obtained by the aid of some
additional reducing agent besides the reducing flame ; thus if the
salt be heated on charcoal instead of a platinum wire, or if a trace
of zinc be added, the result is more quickly obtained. In the
presence of small quantities of iron the bead appears brown-red.

Titanium dioxide is separated from silicon dioxide (or silicates)
by the action of hydrofluoric acid (sulphuric acid being present to
prevent the volatilisation of titanium, as fluoride). Titanium is
obtained in solution (and separated from silica, and also from
compounds of tantalum and niobium) by fusion with hydrogen
potassium sulphate. The " melt " is extracted with cold water,
when the titanium passes into solution as a sulphate.

When titanium oxide is fused with potassium carbonate,
potassium titanate, K 2 TiO 3 , is formed, which, being insoluble in
water, may be separated from any alkaline silicate by extracting
the melt in cold water.

If the residue of potassium titanate be then treated with cold
dilute hydrochloric acid, it dissolves, the solution containing titanic
acid, Ti(HO) 4 or TiO 2 ,2H 2 O. If this solution be heated, the
titanic acid is rendered insoluble, and is therefore precipitated,
the precipitate being the hydrated oxide H 2 TiO 3 , or TiO(HO) 2 ,
or TiO 2 ,H 2 O.

If to the cold solution of titanic acid there be added ammonia,
or the hydroxide or carbonate of the alkalies, or ammonium
sulphide, a white precipitate is produced, consisting of the hydrated
compound Ti(HO) 4 , or TiO 2 ,2H 2 O, which redissolves readily in
dilute hydrochloric acid or sulphuric acid. Moderate rise of
temperature at once converts this soluble compound into the
insoluble titanic acid, H 2 TiO 3 , or TiO 2 ,H 2 O.

Uranium. Like the element chromium, uranium forms
uran07/.r and uram^r salts, as well as urinates.

The uranous salts are derived from uranous oxide, UO 2 , in
which the element is tetravalent. Thus, uranous sulphate is
represented by the formula U(SO 4 )2.

These salts readily pass by oxidation into uranic compounds,
and are therefore powerful reducing agents. They are for the
most part green in colour. The uranic or uranyl salts are derived
from the oxide, UO 3 , or (UCX>)O, in which the element is hexa-
valent. They are regarded as containing the divalent radical
uranyl, (UO 2 ), which takes the place of a divalent metal; thus,
uranyl sulphate and nitrate are expressed by the formulas
(UO 2 )SO 4 ,3H 2 O and (UO 2 )(NO 3 ) 2 ,6H 2 O respectively. These
salts are mostly yellow in colour, and soluble in water.

The uranates are constituted like the dichromates (uranates
corresponding to normal chromates are not known). Thus, sodium
uranate, Na 2 U 2 O 7 , analogous to sodium dichromate, Na 2 Cr 2 O 7 .

Uranyl nitrate and acetate are the salts most commonly met
with, being used in the volumetric determination of phosphoric

Alkaline carbonates give with uranyl salts a yellow precipitate,

74 Qualitative Analysis.

consisting of a double carbonate of uranium and the alkali. The
precipitate is soluble in excess of the reagent.

The uranium in this solution is not precipitated by ammonium
sulphide. If the solution be neutralised with acid, a uranate of
the alkali is thrown down.

With uranous salts, these reagents give a green precipitate,
also soluble in excess.

Caustic alkalies give with uranyl salts a yellow precipitate of
the uranate of the alkali, insoluble in excess of the reagent.

With uranous salts, these reagents give a chocolate-coloured
precipitate of uranous hydroxide, U(HO) 4 .

Ammonium sulphide gives with uranyl salts a brown precipitate
of uranyl sulphide, (UO 2 )S, insoluble in excess of the reagent, but
soluble in normal ammonium carbonate.

In acid solutions sulphuretted hydrogen reduces uranyl to
uranous compounds.

Potassium ferrocyanide gives with iiranyl salts a brown pre-
cipitate. This reagent is employed as an indicator in the volu-
metric estimation of phosphoric acid by means of uranyl salts.



THE metals of this group (as well as those of Group I.) are
characterised by their common property of forming sulphides in an
acid solution ; that is to say, their sulphides are insoluble in dilute
acids. By this property they are all .sharply separated from the
metals of Group III. The acid which has been found to be the
most convenient to have present, is hydrochloric acid ; and since,
on acidifying with this acid preparatory to the precipitation of the
sulphides, the chlorides of silver, lead, and mercicrous mercury * are

* Mercury forms two classes of soils,' mercuric and mercurous ; and the
relation in which they stand to each other is very interesting. In the com-
pounds of the first type, the metal is playing the part of an ordinary divalent
element, replacing two atoms of hydrogen in acids and giving such salts as
HgCl 2 , Hg(NO 3 ) 2 , HgSO 4 , HgS, etc. In the mercurous compounds the pro-
portion of mercury to the negative radical is twice as great as in the mercuric
salts, hence the composition of, say, mercurous chloride may be expressed
either by the formula Hg 2 Cl 2 or HgCl. Some chemists adopt the latter
formula, and, regarding the mercury in mercurous compounds as acting the
part of a monovalent element, express the various salts by such formulae as
HgXO 3 , Hg 2 SO 4 , etc. Others prefer to consider the mercurous salts as com-
pounds, in which the divalent radical, or double atom, (Hg 2 ), is substituted for
the single divalent atom, (Hg), and therefore express the compounds by such
formulas as Hg 2 Cl 2 , Hg 2 (NO 3 ) 2 , Hg 2 SO 4 . The density of the vapour yielded
by heating mercurous chloride is 117 '59, which, being half that demanded by
the formula Hg 2 Cl 2 , gave support to the view that HgCl was the correct
formula. But it has since been shown that the compound dissociates on
heating, into mercuric chloride, HgCl 2 , and Hg (mercury giving monatomic
molecules). In the absence of any conclusive evidence in favour of either view,
the formulae adopted for mercurous compounds in this book will be those which
represent them as containing the double atom (Hg 2 ). As the two classes of
compounds present great differences in their chemical reactions, behaving,
indeed, more like compounds of two different metals, one of which belongs to
Group I. and the other to Group II., there is at least some advantage in
employing a symbol for mercurous mercury (Hg 2 ), which at a glance dis-
tinguishes it from that used to denote mercuric mercury, Hg. All the salts of
mercury will therefore be formulated as salts of a dibasic metallic radical,
which in the mercuric compounds has the symbol Hg, and in the mercurous
salts (Hg 2 ) (with or without the bracket). Thus

Chlorides, HgCl 2 (Hg,)Cl 2

Sulphates, HgS0 4 (Hg 2 )SO 4

76 Qualitative Analysis.

thrown down, these three metals are treated separately, and con-
stitute Group I.

The precipitation of these three metals as chlorides, however,
is only complete in the case of Ag and (Hg 2 ). Lead chloride is
soluble to some extent even in cold water, hence a portion of the
lead passes through into Group II.

The metals of Group II. are divided into two sections, namely

(1) Metals whose sulphides are insoluble in ammonium

Mercury (mercuric), lead, bismuth, cadmium, copper.

(2) Metals whose sulphides dissolve in ammonium sulphide

Arsenic, antimony, tin, gold, platinum.

Mercury, Hg.

DRY REACTiONS.-^When heated alone in a tube, many mercury
compounds (those with the halogens, for example) volatilise un-
changed, giving sublimates of the same compound. The iodide
(red) when heated forms a sublimate, consisting chiefly of the yellow
allotropic form of HgI 2 , which when cold changes to red if scratched
or rubbed. Some mercury compounds when heated decompose,
and metallic mercury volatilises and sublimes in the tube.

If a mercury salt be mixed with several times its weight of
sodium carbonate (both being as dry as possible), and the mixture
be strongly heated in a dry narrow test-tube, a sublimate of metallic
mercury will be obtained. The sublimed mercury will present the
appearance of a bright metallic mirror, but if examined by means
of a lens, or if rubbed with a glass rod, distinct globules of liquid
metal will be visible.

WET REACTIONS. (a) Mercuric Compounds. Of the com-
mon salts, the nitrate, sulphate, chloride, and bromide (but not the
iodide) are soluble in water, but the solubility is not very great.

KHO or NaHO gives with mercuric compounds a yellow
precipitate * of mercuric oxide, HgO

HgCl 2 + 2KHO = HgO + H 2 O + 2KC1

Nitrates, Hg(NO 3 ) (Hg 2 )(NO 3 ).>

Basic nitrates, Hg(N0 3 ) 2 ,2HgO,H 2 ... iHg,)(NO,),,{Hg a )O,H t O
Double ammonium compounds, NH 2 HgCl NH 2 (Hg 2 )Cl

* On the first addition of the reagent, the precipitate appears a brownish
colour (probably due to the momentary formation of the hydroxide, which is
incapable of existing), but almost immediately it becomes yellow. Why the
oxide obtained by precipitation should be yellow, while that prepared in the
dry way is brick-red, is not known. Compare also the sulphide.

Group II. Division I. 77

The precipitate is insoluble in excess of the reagent.

NH 4 HO produces a white precipitate of an ammoniacal mercuric
compound, where two atoms of hydrogen from the ammonium
radical are replaced by the divalent atom Hg ; thus

HgCl 2 + 2 NH 4 HO = NH.HgCl + NH 4 C1 + 2H 2 O

i " ^^^^ -^

Or with mercuric nitrate

Hg(N0 3 ) 2 + 2NH 4 HO = (NH 2 Hg)N0 3 + NH 4 NO 3 + 2H 2 O

H 2 S produces a black * precipitate of HgS. The precipitation
is only complete after some time, and when the solution is con-
siderably dilute. The compound is insoluble in HC1, and in
HNO 3 even when boiling. (The prolonged action of boiling HNO 3
partially converts it into the white compound Hg(NO 3 ) 2 , 2HgS.)
Mercuric sulphide dissolves in aqua regia, forming mercuric
chloride. In the presence of caustic alkalies it dissolves in sodium
or potassium sulphide (not in ammonium sulphide], forming the
double sulphides, HgS,Na 2 S and HgS,K 2 S.

(NH 4 ),S gives the same precipitate. The same result is also
obtained by the addition of sodium thiosulphate, Na 2 S 2 O 3 , to a warm
solution acidified with HC1.

II precipitates HgI 2 as a rich scarlet compound, soluble in
excess of either solution. When first precipitated it appears
yellow, but quickly turns salmon-red and then scarlet. The com-
pound is dimorphous, and can be obtained either in the red
(quadratic crystals) or the yellow (rhombic prisms'] variety.

Reduction of Mercuric Compounds. By reducing agents
mercuric compounds may be converted into mercurous salts, or the
reduction may go a stage further and result in the precipitation of
mercury in the metallic state. Thus, on the addition of stannous

* The progress of this precipitation is accompanied by very characteristic
changes of colour. The first action of the H 2 S is to give a -white precipitate,
which then passes through various shades of colour, from yellow to yellowish-
red, to brown, and lastly, black. The white substance is a compound of HgS
with the mercuric salt in solution, HgCl 2 ,2HgS, or Hg(NO ? )2 ,2HgS, and the
changes in colour are ascribed to the gradual conversion of this into black, HgS.
It does not appear quite obvious, however, how a gradual alteration in the pro-
portions of the white double compound and the black sulphide can give the
orange and reddish tints which are seen. Mercuric sulphide prepared by other
processes is red (the pigment known as vermilion}. Why the compounds pre-
pared in different ways should be so very different in colour is not known ;
probably it is a case of dimorphism similar to that exhibited by HgI 2 , and it
may be that to some extent the red HgS is precipitated, but is not stable
under the conditions which are present, t and so passes into the black

78 Qualitative Analysis.

chloride, SnCl 2 , a white precipitate of mercurous chloride is

2 HgCl 2 + SnCl 2 = (Hg 2 )Cl 2 + SnCl 4

On gently warming with an excess of stannous chloride, the pre-
cipitated mercurous chloride changes to a grey deposit of mercury
in a condition of fine powder

(Hg) 2 Cl 2 + SnCl 2 = SnCl 4 + 2Hg

Many metals are capable of displacing mercury from its solutions,
the mercury being deposited upon the metal. Thus, if a strip of
clean copper be immersed in a neutral or slightly acid solution of a
mercury salt, it becomes coated with a white silvery deposit of an
amalgam of copper and mercury, from which the mercury can be
readily volatilised and obtained as a metallic sublimate by heating
the copper in a dry test-tube. In the case of mercuric salts, the
action may be regarded as taking place in two stages ; thus

2Hg(N0 3 ) 2 + Cu = (Hg 2 )(N0 3 ) 2 + Cu(N0 3 ) 2
(Hg 2 )(N0 3 ) 2 + Cu = 2Hg + Cu(N0 3 ) 2

(fi) Mercurous Compounds. Of the common salts mer-
curous nitrate is the only one which is readily soluble, and this only
so long as the water is acid with nitric acid. The addition of much
water results in the precipitation of a basic nitrate. Mercurous
sulphate is soluble with difficulty.

KHO or NaHO throws down a black precipitate of (Hg 2 )O.
Mercurous oxide is very unstable. When gently warmed, or even
upon exposure to light, it is converted into HgO and Hg.

NH HO precipitates an ammoniacal mercurous compound,
which is black. Its composition is exactly analogous to the
corresponding mercuric compound

Hg 2 (N0 3 ) 2 + 2NH 4 HO = NH 2 (Hg 2 )N0 3 + NH 4 NO 3
H 2 S produces a black precipitate, which is a mixture of HgS
and Hg. (Hg 2 S is not known to exist.) This precipitate, therefore,
behaves, on treatment with nitric acid, in the same way as that
obtained from a mercuric solution. If we imagine the free atom
of mercury as first dissolving in the acid, the mercuric nitrate so
formed unites with HgS, giving the white insoluble compound
Hg(N0 3 ) 2 ,2HgS.

(NH 4 ) 2 S gives the same precipitate, but in this case the free
mercury will be also converted into HgS in proportion as the
ammonium sulphide contains more or less polysulphide, for alkaline
poly sulphides convert metallic mercury into HgS.

Group II. Division i. 79

HC1 and soluble chlorides precipitate white mercurous chloride,
Hg, 2 Cl 2 . Insoluble in dilute acids ; soluble in boiling HNO 3 , being
converted into HgCl 2 and Hg, and the mercury then dissolves to
mercuric nitrate, with evolution of oxides of nitrogen.

Long boiling with concentrated HC1 decomposes Hg 2 Cl 2 into
HgCl 2 (which dissolves) and Hg, which separates (mercury being
insoluble in HC1). Chlorine water converts it into mercuric
chloride ; thus

Hg 2 Cl 2 + C1 2 = 2HgCl 2

Ammonia converts it into black mercurous ammonium chloride,
NH 2 (Hg 2 )Cl. (This constitutes one of the most characteristic
reactions for mercurous compounds.)

Mercurous salts are reduced to metallic mercury by the reducing
agents which reduce the mercuric compounds ; thus, with stannous
chloride a grey precipitate of mercury is at once produced

Hg 2 (NO 3 ) 2 + SnCl 2 + 2HC1 = SnCl 4 = 2HNO 3 + 2 Hg

Lead, Pb.

DRY REACTIONS. Lead compounds are very readily reduced
when heated upon charcoal before the blowpipe flame, either alone
or mixed with sodium carbonate or potassium cyanide. Globules
of metallic lead are thus obtained, and at the same time a yellowish
incrustation is formed, consisting of the oxide, PbO (litharge).
When cold, one of the globules can be removed and the properties
of the metal examined. Lead may be recognised by its malleability
and softness, the latter property enabling it to leave a black mark
when rubbed upon paper. It is insoluble in cold HC1 or H 2 SO 4 ,
but readily dissolves in HNO 3 , forming Pb(NO 3 ) 2 , which, being
insoluble in nitric acid, remains as a white deposit, but which
dissolves on dilution with water.

WET REACTIONS. The only salts of lead which are met with
in analysis are derived from plumbic oxide, PbO, in which the
metal is divalent.* Of the common salts, the nitrate and acetate
are readily soluble in water ; the chloride, bromide, and iodide are
sparingly soluble.

KHO, NaHO, or NH 4 HO gives a white precipitate of either
lead hydroxide or a basic compound, depending upon whether the
lead solution or the alkali is in excess all the time. For example,
if the potash be added to the lead solution, the precipitation

* Salts are known in which lead is tetravalent, e.g. lead tetracetate,
Pb(GtHiO*)4. The compound PbCl 4 has also b:en obtained.

82 Qualitative Analysis.

In the case of bismuth chloride, the oxychloride is thrown

BiCl 3 + H 2 O = BiOCl + 2HC1

This compound is not so easily dissolved by HC1 as the basic
nitrate is by HNO 3 , therefore the reaction is complete, the whole
of the bismuth being precipitated if the solution is dilute.

KHO, NaHO, or NH 4 HO precipitates the white hydroxide
Bi(HO) 3 , or Bi 2 O 3 ,3H 2 O.* Insoluble in excess of the precipitants.
From boiling solutions, or on heating to boiling, the monohydrated
oxide is formed, Bi 2 O 3 ,H 2 O.

K 2 CO 3 , Na 2 CO 3 , or (NH 4 ) 2 CO 3 throws down a white basic
carbonate, (BiO) 2 CO 3 .f Insoluble in excess of the reagents.

H 2 S or (NH 4 ) 2 S precipitates bismuthous sulphide, Bi 2 S 3 , as a
dark brown, almost black, compound. Soluble in HNO 3 ; insoluble
in alkaline sulphides.^

Sulphuric acid produces no precipitate with a bismuth salt.
Potassium dichromate, K 2 Cr 2 O 7 , precipitates basic bismuth dichro-
mate, (BiO) 2 Cr 2 O 7 . Insoluble in KHO. (These two reactions
distinguish between Pb and Bi.)

TheNmost characteristic reaction for bismuth is the formation
of the insoluble oxychloride, BiOCl, on the addition of water to an
acid solution of BiCl 3 . As explained above, the reaction with the
chloride is more delicate than with the nitrate, hence, if the nitrate
is used, it should be converted into the oxychloride. This can be

Online LibraryGeorge Samuel NewthA manual of chemical analysis, qualitative and quantitative → online text (page 8 of 45)