A. B. (Albert Benjamin) Prescott.

Qualitative chemical analysis; a guide in qualitative work, with data for analytical operations and laboratory methods online

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ignited (Cleve, J., 1874, 261). The hydroxide, Th(OH) 4 , is formed by precipita-
tion of the salts by the alkalis. It is a white, heavy, gelatinous precipitate,
drying to a hard glassy mass. The chloride, ThCl 4 , and the nitrate, Th(N0 3 ) 4 ,
*>re deliquescent. The chloride is a white body melting at a white heat and then
subliming in beautiful white needles (Kruess and Nilson, I.e.). The sulphate
is soluble in five parts of cold watiT. The carbon-file, oxalate and phosphate are
insoluble in water; the a.ralatc is scarcely soluble in dilute mineral acids.
Alkali hydroxides or sulphides precipitate thorium hydroxide, Th(OH) 4 ,
insoluble in excess of the reagent. Tartaric and citric acids hinder the pre-
cipitation. Alkali carbonates precipitate the basic carbonate, soluble in ex-
cess, if the reagent be concentrated. The solution in (NH 4 ) 2 C0 3 readily repre-
cipitates upon warming. BaC0 3 precipitates thorium salts completely/ Oxalic
acid and oxalates form a white precipitate (distinction from Al and Gl), not
soluble in oxalic acid or in dilute mineral acids; soluble in hot concentrated
(NH 1 ),C,O 4 and not reprecipitated on cooling and diluting (distinction from
Ce and La). A saturated solution of K 2 S0 4 slowly but completely precipitates
a solution of Th('S0 4 ) 2 , forming potassium thorium sulphate; insoluble in a
saturated K,SO 4 solution, sparingly soluble in cold water, readily soluble in
hot water. HF precipitates Th.F 4 , insoluble in excess, gelatinous, becoming
crystalline on standing. Boiling freshly precipitated Th(OH) 4 with KF in
presence of HF forms K 2 ThF 6 .4H 2 O , a heavy fine white precipitate almost
insoluble in water. The distinguishing reactions of thorium are the precipitation
with oxalates and with K 2 SO 4 , and failure to form a soluble compound on
fusion with Na 2 C0 3 (distinction from Si0 2 and Ti0 2 ).

170. Titanium. Ti = 48.1. Valence three and four.

Titanium is found quite widely disiribut3d as lutiie, Lrookite, anatase,
titanite, titaniferous iron, FeTiOs, and in many soils and clays. Never found
native. It is prepared by heating the fluoride or chloride with K or Na . It
is a dark gray powder, which shows distinctly metallic when magnified', melt-
ing point, 1800 (Cir. B. of S., 1915). Heated in the air it burns with an unusu-
ally brilliant incandescence', sifted into the flame it burns with a blinding bril-
liance. Chlorine in the cold is without action, when heated it combines with
vivid incandescence. It decomposes water at 100. It is soluble in acids, with
evolution of hydrogen, forming titanous salts. At a higher temperature it com-
bines directly with Br and I . It is almost the only metal that combines
directly with nitrogen when heated in the air (Woehler and Deville, A.,
1857, 103, 230; Merz, J. pr., 1866, 99, 157). The most common oxide of
titanium is the dioxide, TiO 2 , analogous to CO 2 and SiO 2 . It occurs more
or less pure in nature as rutile, brookite and anatase; it is formed by igni-
tion of the hydrated titanic acid or of ammonium titanate (Woehler, J., 1849,
268). Ignition of TiO 2 in dry hydrogen gives Ti 2 O 2 , an amorphous black
powder, dissolving in H 2 SO 4 to a violet-colored solution (Ebelmen, A. Ch.,
1847, (3), 20, 392). TiO is formed when TiO 2 is ignited with Mg:2TiO 2 -f-
Mg = TiO + MgTiO 3 (Winkler, B., 1890, 23, 2660). Other oxides have been
reported. Titanic acid, TiO 2 , is a white powder, melts somewhat easier than
SiO 2 , soluble in the alkalis unless previously strongly ignited. Mixed with
charcoal and heated in a current of chlorine, TiCl 4 is formed. The bromide
is formed in a similar manner. TiO 2 acts as a base, forming a series of stable
salts; also as an acid, forming titanates. TiCli is a colorless liquid, fuming
in the air; it boils at 136.41 (Thorpe, J. C., 1880, 37, 329); it is decomposed
by water, forming titanic acid, which remains in solution in the HC1 present.
Solutions of most of the titanic salts, when boiled, deposit the insoluble
meta-titanic acid. HF dissolves all forms^ of titanic acid; if the solution
be evaporated in presence of H 2 SO 4 no TiF., is volatilized (distinction from
SiF 4 ). When evaporated with HF alone, TiF 4 is volatilized. The double
potassium titanium fluoride, K 2 TiF 6 , formed by fusing TiO 2 with acid KF , is
sparingly soluble in water (96 parts), readily soluble in HC1 . Solutions of
titanic salts in water or acid solutions of titanic acid are precipitated by
alkali hydroxides, carbonates and sulphides as the hydrated titanic acid, insoluble
in excess of the precipitants and in ammonium salts. BaCO 3 gives the same

206 URANIUM. 171 -

precipitate. K 4 Fe(CN) gives a reddish-yellow precipitate; K,Fe(CN) 6 a yellow
precipitate. NaaHPO, precipitates the titanium ahinjxt cviuplctcliL even in the
presence of strong- HC1 . An acid solution of TiO, when treated with Sn or
Zn gives a pale blue to violet coloration to the solution, due to a partial reduction
of the titanium to the triad condition. These colored solutions are precipitated
by alkali hydroxides, carbonates and sulphides. ITS is without action. The
solution reduces Fe"' to Fe" , CM" to Cu' , and salts of Hg- , Ag and Au to the
metallic state; the titanium becoming again the tetrad. The reduction by Sn
or Zn takes place in presence of HF (distinction from columbic acid). Titanium
compounds fused in the flame with microcosmic salt give in the reducing flame
a yellow bead when hot, cooling to reddish and violet (reduction of the tita-
nium). With FeSO t in the reducing flame a Mood-red bead is obtained.

Titanium is very readily detected in minerals as follows. 0.1 gram of the
finely powdered mineral is mixed with 0.2 gram of finely powdered sodium
fluoride and 3 grams sodium pyrosulphate added without mixing. The crucible
is heated until copious sulphuric acid fumes are evolved. The fused mass is
rapidly cooled and heated with 2-3 c.c. dilute sulphuric acid and 10 c.c. water
added. The solution is dividqd into two parts and a few drops of hydrogen
peroxide added to one part. A yellow color is produced by the titanium.
Chlorides, bromides and iodides interfere with this very delicate reaction (Weber,
Z., 40, 799, Noyes, J. Soc. Ind., 10, 485).

171. Uranium, U = 238.2. Valence four and six.

Specific gravity, 18.685 (Zimmermann, A., 1882, 213, 285). Melting point,
<1850 (Cir. B. of S., 1915). Found in various minerals; its chief ore is
pitch-blende, which contains from 40 to 90 per cent of UsOg . Prepared by
fusing UClt with K or Na (Zimmermann, A., 1883, 216, 1; 1886, 232, 273).
It has the color of nickel, hard, but softer than steel, malleable, permanent
in the air and water at ordinary temperatures; when ignited burns with incan-
descence to UgOg J unites directly with Cl , Br , I and S when heated; soluble
in HC1 , HoSO 4 and slowly in HNO 3 . Uranous oxide, UO2 , formed by ignit-
ing the higher oxides in carbon or hydrogen, is a brown powder, soon turning
yellow by absorption of oxygen from the air. Uranous hydroxide is formed
by precipitating uranous salts with alkalis. Uranic oxide, UOs , is formed
by heating uranic nitrate cautiously to 25, and upon ignition in the air both
this and other uranium oxides, hydroxides and uranium oxysalts with volatile
acids are converted into U 3 O 8 = UO 2 2UO 3 . Uranium acts as a base in two
classes of salts, uranous and uranyl salts. Uranous salts are green and give
green solutions, from w y hich alkalis precipitate uranous hydroxide, insoluble in
excess of the alkali; alkali carbonates precipitate U(OH) 4 , soluble in
(NH 4 ) 2 CO 3 ; with BaCO 3 the precipitation is complete even in the cold. H 2 S is
without action; (NH 4 )J3 gives a dark-brown precipitate; K 4 Fe(CN) 6 gives a
reddish-brown precipitate. In their action toward oxidizing and reducing
agents uranous and uranyl (uranic) salts resemble closely ferrous and ferric
salts; uranous salts are even more easily oxidized than ferrous salts, e. g., by
exposure to the air, by HNO 3 , Cl , HC1O 3 , Br , KMnO 4 , etc. Gold, silver and
platinum salts are reduced to the free metal. The hexad uranium (U^ 1 ) acts
as a base, but usually forms basic salts, never normal: we have TJ0 2 (NO :J ) 2 ,
not TT(NO 3 ) ; UO L S0 4 , not TT($O 4 ) 3 . These basic salts w T ere formerly called
uranic salts, but at present (ir0 2 )" is regarded as a basic radical and called
tirunjil, and its salts are called uranyl salts, e.g., UO 2 C1 2 uranyl chloride,
(TTO 2 ) S (PO 4 ) 2 uranyl orthophosphate. Solutions of uranyl salts are yellow;
KOH and NaOH give a yellow precipitate, uranates, K 2 TJ 2 7 and Na.ILO, ,
insoluble in excess. Alkali carbonates give a yellow precipitate, soluble in
excess; BaC0 3 and CaC0 3 give TJ0 3 . H 2 S does not precipitate the uranium,
but slowly reduces uranyl salts to uranous salts (Formanek, A., 1890, 257, 115)1
(NH 4 )..S gives a dark-brown precipitate. K 4 Fe(CN) gives a reddish-brown
precipitate. Used in the analysis and separation of uranium compounds
(Fresenius and Hintz, Z. angeic., 1895, 502). Sodium phosphate gives a yellow
precipitate. The hexad uranium acts as an acid toward some stronger bases.

171 rt, 8. VANADIUM- 207

Thus we have K 2 TT2O 7 and Na^U^Oy , formed by precipitating uranyl salts with KOH
and NaOH ; compare the similar salts of the hcxad chromium, K 2 Cr 2 O 7 and
Na2Cr 2 O 7 . Other oxides of uranium are described, but are doubtless combinations
of UO 2 and UO 3 . Zn , Cd , Sn , Pb , Co , Cu , Fe , and ferrous salts reduce uranyl
salts to uranous salts. Solutions of Sn, Ft, Au, Cu, Hg and Ag are reduced to
the metal by metallic uranium (Zimmcrmann, I.e.). For method of recovery of
waste uranium compounds, see Laube (Z. angew., 1889, 575).

171a. Vanadium. V = 51.0. Valence two to five.

1. Properties. -Specie gravity, 5.8 (Moissan, C. r., 122); melting point 1720
(Cir. B. S., 35, 1915). A grayish non-magnetic powder; slowly oxidized in the air,
rapidly on ignition with formation of V 2 O 5 . It forms with chlorine the dark
brown tetrachloride.

2. Occurrence. It is often found in iron and copper ores and in some clays and
rare minerals, e.g., vanadinite, (Pb 5 Cl(VO 4 )3) ; volborthite, (Cu,Ca,Ba) 3 (OH) 3 VO4
+6H 2 O) ; mottramite (a hydrous vanadate of lead, copper, and other
divalent elements, of uncertain formula, (R 3 (VO 4 ) 2 .3R(OH) 2 )) .

3. Preparation. The vanadium ores are treated chiefly for the preparation of
ammonium vanadate and vanadic acid. The ores are fused with KN0 3 , form-
ing- potassium vanadate. This is precipitated with Pb or Ba salts and then
decomposed with H.SO 4 . The vanadic acid is neutralized with NH,OH and
precipitated with NH 4 C1 , in which it is insoluble. This upon ignition gives
V,O pure (\Yohler, A., 1851, 78, 125). The metal is prepared from the dichlo-
ride, VCL , by long-continued ignition in a current of hydrogen.

4. Oxides. Vanadium forms four oxides: VO , gray; V 2 3 , black; V0 2 , dark
blue; and V,O,, , dark red to orange red.

5. Solubilities. Vanadium is not attacked by dilute HC1 or H 2 SO 4 ; concen-
trated H,SO 4 gives a greenish-yellow solution; HNO 3 a blue solution. VO dis-
solves in acids to a blue solution with evolution of hydrogen. V 2 O 3 dissolves
in dilute HC1 to a dark greenish-black solution. Chlorine forms with V 2 O 3 ,
VOC1 3 and V 2 O r , . VO., dissolves in acids to a blue solution, from which solu-
tions Na.CO, gives a precipitate of V 2 2 (OH) 4 + 5H 2 O , grayish-white mass,
losing 4H 2 O at 100 and turning black, soluble in acids and alkalis. V 2 O 5
exists in several modifications with different solubilities in water, the red
modification being- soluble in 125 parts of water at 20 (Ditte, C. r., 1880, 101,
698). Vanadic acid forms three series of salts, ortho, meta and pyro, analogous
to the phosphates. Most salts are the metavanadates. The ortho compounds
are quite unstable, readily changed to the meta and pyro compounds. Alkali
vanaclates are soluble in water, the ammonium vanadate least soluble and not
at all in NH 4 C1 .

6. Reactions. Solutions of vanadic acid produce brown precipitates with
alkalis, soluble in excess to a yellowish-brown color. Potassium ferrocyanide
gives a green precipitate, insoluble in acids. Tannic acid gives a blue-black
solution, which is said to make a desirable ink. Ammonium sulphide precipi-
tates V,S 5 , brown, soluble with some difficulty in excess of the reagent to a
reddish-brown thio salt. From this solution acids reprecipitate the brown
vanadic sulphide, V 2 S 5 .

If to a solution of a vanadate, neutral or alkaline, solid NH 4 C1 be added, the
vanadium is completely precipitated as NH 4 V0 3 , ammonium metavanadate,
crystalline, colorless, insoluble in NH 4 C1 solution; upon ignition in air or oxy-
gen, pure vanadic oxide, V 2 B , is obtained.

7. Ignition. Borax gives with vanadium compounds in the outer flame a
colorless bead, yellow if much vanadium be present; in the inner flame a green
bead, or brow r n when vanadium is present in large quantities and hot, becoming
green upon cooling. All the lower oxides of vanadium ignited in air or
oxygen give V 2 O 5 .

8. Detection. Vanadium will almost always be found as a vanadate (2) and
is detected by the reactions used in its purification (3) ; also by the reactions
with reducing agents, forming- the colored lower oxidized compounds (10).

208 YTTERBIUMYTTRIUM. 171 rt, 10.

9. Estimation. (1} It is precipitated as basic lead vanadate and dried at
100. (2} It is precipitated as ammonium vanadate, NH 4 VO 3 , in strong
NH 4 C1 solution, ignited to the oxide V-Oa , and weighed.

10. Oxidation. Zn , in solutions of vanadates with dilute H 2 SOi , reduces the
vanadium to the tetrad, a green to blue solution, then greenish-blue to green,
the triad, and finally to lavender blue, the dyad. H 2 S reduces vanadates to the
tetrad with separation of sulphur. Oxalic acid and sulphurous acid also reduce
vanadates to the tetrad, the solution becoming blue.

172. Ytterbium. Yb = 173.5. Valence three.

Obtained as an earth by Marignac (C. r.. 1878, 87, 578) from a gadolinite
earth; by Delafontaine (C. r., 1878, 87, 933) from sipylite found at Amherst, Va.
Nilson (/?., 1879, 12, 550; 1880, 13, 1433) describes its preparation from euxenite
and its separation from Sc . It has the lowest bacisity of the yttrium earths.
The double potassium ytterbium sulphate is easily soluble in water and in
potassium sulphate. The oxalate forms a white crystalline precipitate, in-
soluble in water and in dilute acids. The salts are colorless and give no
absorption spectrum. For the spark spectrum see Welsbach (J/., 1884, 5, 1).
The oxide, Yb,,0 3 , is a white powder, slowly soluble in cold acids, readily upon
warming. The Jiydro.riilc forms a gelatinous precipitate, insoluble in i^H 4 OH
but soluble in KOH . It absorbs CO, from the air. The nitrate melts in it
water of crystallization and is very soluble in water.

173. Yttrium. Y = 88.7. Valence three.

Yttrium is one of the numerous rare metals found in the gadolinite mineral
at Ytterby, near Stockholm, Sweden: also found in Colorado (Hidden and
Mackintosh, Am-. '., 1889, 38, 474). The metal has been prepared by electro-
lysis of the chloride; also by heating the oxide, Y 2 3 , with Mg- (Winkler, B.,
1890, 23, 787). Melting point, 1490 (Cir. B. of S.~ 1915). The study of these
rare earths is by no means complete. It is also claimed that they have not
yet been obtained pure, but that the so-called pure oxides really consist of a
mixture of oxides of from five to twenty elements (Crookes, C. N., 1887, 55,
107, 119 and 131). The most of these rare earths do not give an absorption
spectrum, but give characteristic spark spectra; and it is largely by this means
that the supposedly pure oxides have been shown to be mixtures of the oxides
of several closely related elements (Welsbach, M., 1883, 4, 641; Dennis and
Chamot, J. Am. Soc., 1897, 19, 799). Yttrium salts are precipitated by the
ai -alis and by the alkali sulphides as the hydroxide, Y(OH) ? , a white bulky pre-
cipitate, insoluble in the excess of the reagents (distinction from Gl). The
oxide and hydroxide are readily soluble in acids; boiling with NH,C1 causes
solution of the hydroxide as the chloride. The alkali carbonates precipitate
the carbonate Y 2 (CO 3 )3 , soluble in a large excess of the reagents. If the solu-
tion in ammonium carbonate be boiled, the hydroxide is precipitated. Soluble
oxalates precipitate yttrium salts as the white oxalate (distinction from Al
and Gl); soluble with some difficulty in HC1 . The double sulphate with
potassium is soluble in water and in potassium sulphate (distinction from
thorium, zirconium and the cerite metals). BaCO 3 forms no precipitate in
the cold (distinction from Al , Fe'" , Cr'" , Th , Ce , La , Nd and Pr). Hydro-
fluoric acid precipitates the gelatinous fluoride, YF 3 , insoluble in water and
in HF . The precipitation of yttrium salts is not hindered by the presence
of tartaric acid (distinction from Al , Gl , Th and Zr). The analysis of yttrium
usually consists in its detection and separation in gadolinite (silicate of Y ,
Gl , Fe , Mn , Ce and La). Fuse with alkali carbonate, decompose with KC1 ,
and filter from the SiO 2 . Neutralize the filtrate and precipitate the Y , La
and Ce as oxalates with (NH 4 ) 2 C 2 O 1 . Ignite the precipitate and dissolve in
HC1 . Precipitate the La and Ce as the double potassium sulphates, and from
the filtrate precipitate the yttrium as the hydroxide with NH,OH . Ignite
and weigh as the oxide. In order to effect complete separations the operations
should be repeated several times.


174. Zirconium. Zr = 90. G. Valence four.

Zirconium is a rare metal found in various minerals, chiefly in zircon,
a silicate; never found native. The metal was first prepared by Berzelius
in 1824 by fusion of the potassium zirconium fluoride with potassium (Pogg.,
1825, 4, 117). Also prepared by electrolysis of the chloride (Becquerel, A.
CA, 1831, 48, 337). Melting point, 1700 ? (Cir. B. of S., 1915). The metal
exists in three modifications: crystalline, graphitoidal and amorphous. The
amorphous zirconium is a velvet-black powder, burning when heated in the
air. Acids attack it slowly even when hot, except HF , which dissolves it in
the cold. It forms but one oxide, ZrO 2 , analogous to SiO 2 and TiO> . ZrO 2
is prepared from the mineral zircon by fusion with a fixed alkali. Digestion
in water removes the most of the silicate, leaving the alkali zirconate as a
sandy powder. Digestion with HC1 precipitates the last of the SiOa and
dissolves the zirconate. The solution is neutralized, strongly diluted and
boiled, whereupon the zirconium precipitates as the basic chloride free from
iron. Or the zirconium may be precipitated by a saturated solution of K 2 SO 4 ,
and after resolution in acids precipitated by NH 4 OH and ignited to ZrO_>
(Berlin, J. pr., 1853, 58, 145; Roerdam, C. C., 1889, 533). ZrO 2 is a white
infusible powder, giving out an intense white light when heated; it shows no
lines in the spectrum. It is much used with other rare earths, La.oO ;( , Y 2 O 3 ,
etc., to form the mantles used in the Welsbach gas-burners (Drossbach, C. C.,
1891, 772; Welsbach, </., 1887, 2670; C. N., 1887, 55, 192). The oxide (or hydroxide
precipitated hot) dissolves with difficulty in acids to form salts. The hydroxide,
ZrO(OH) 2 , precipitated in the cold dissolves readily in acids. As an acid,
zirconium hydroxide, ZrO(OH) 2 = H 2 ZrO 3 , forms zirconates, decomposed by
acids. As a base it forms zirconium salts with acids. The sulphate is easily
soluble in water, crystallizing from solution with 4H.O . The phosphate is
insoluble in water, formed by precipitation of zirconium salts by Na 2 HP0 4 or
H 3 PO 4 . The silicate, ZrO 2 .SiO 2 , is found in nature as the mineral zircon,
usually containing traces of iron. Zirconium chloride is formed when a current
of chlorine is passed over heated ZrO 2 . mixed with charcoal. It is a white
solid, may be sublimed, is soluble in water. Solutions of zirconium salts are
precipitated as the hydroxide, ZrO(OH) 2 , by alkali hydroxides and sulphides,
a white flocculent precipitate, insoluble in excess of the reagents, insoluble in
NH 4 C1 solution (difference from Gl). Tartaric acid prevents the precipitation.
Alkali carbonates precipitate basic zirconium carbonate, white, soluble in
excess of KHCO 3 or (NH 4 ) 2 C0 3 ; boiling precipitates a gelatinous hydroxide
from the latter solution. BaCO 3 does not precipitate zirconium salts com-
pletely, even on boiling. The precipitates of the hydroxide and carbonate are
soluble in acids. Oxalic acid and oxalates precipitate zirconium oxalate, solu-
ble in excess of oxalic acid on warming, and soluble in the cold in (NH 4 ) 2 C 2 4
(difference from thorium); soluble in HC1 . A saturated solution of K 2 S0 4
precipitates the double potassium zirconium sulphate, white, insoluble in excess
of the reagent if precipitated cold, soluble in excess of HC1; if precipitated
hot, almost absolutely insoluble in water or HC1 (distinction from Th and Ce).
Zirconium salts are precipitated on warming with Na,S.,O 3 (separation from
Y, Nd and Pr). Solution of H,O, completely precipitates zirconium salts.
Tumcric paper moistened with a solution of zirconiiim salt and HC1 is colored
orange upon drying (boric acid gives the same reaction) (Brush, J. pr., 1854,
62, 7). HF does not precipitate zirconium solutions, as zirconium fluoride,
ZrF 4 , is soluble in water and in HF (distinction from Th and Y).


Barium. Ba = 137.37. Calcium. Ca = 40.07.

Strontium. Sr = 87.03. Magnesium. Mg = 24.32.

175. Like the alkali metals, Ba , Sr , and Ca oxidize rapidly in the ai-

at ordinary temperatures forming alkaline earths and dcnniijtone wuft >\

forming hydroxides with evolution of heat. Mg oxidizes rapidly in the air


when ignited, decomposes water at 100, and its oxide in physical proper-
ties farther removed from Ba , Sr , and Ca than these oxides are from e;icli
other slowly unites with water without sensible production of heat. As
compounds, these metals are not easily oxidized beyond their quantivalence
as dyads, and they require very strong reducing agents to restore them
to the elemental state.

176. In basic power, Ba is the strongest of the four, Sr somewhat
stronger than Ca, and Mg much weaker than the other three. It will be
observed that the solubility of their hydroxides varies in the same decreas-
ing gradation, which is also that of their atomic weights; while the
solubility of their sulphates varies in a reverse order, as follows: (7) :

177. The hydroxide of Ba dissolves in about 30 parts of water; that of
Sr, in 100 parts; of Ca, in 800 parts; and of Mg, in 100,000 parts. The
sulphate of Ba is not appreciably soluble in water (429,700 parts at 18.4;
Hollemann, Z. phys. Ch., 1893, 12, 131); that of Sr dissolves in 10,000
parts; of Ca , in 500 parts; of Mg , in 3 parts. To the extent in which they
dissolve in water, alkaline earths render their solutions caustic to the
taste and touch, and alkaline to test-papers and phenolphthalein.

178. The carbonates of the alkaline earths are not entirely insoluble
in pure water: BaC0 3 is soluble in 45,566 parts at 24.2 (Hollemann,
Zeit. phys. Ch., 1893, 12, 125); SrC0 3 in 90,909 parts at 18 (Kohlrausch
and Rose, Zeit. phys. Ch., 1893, 12, 241); CaC0 3 in 80,040 parts at 23.8
(Hollemann, /. c.); MgC0 3 in 9,434 parts (Chevalet, Z., 1869, 8, 91). The
presence of NH 4 OH and (NH 4 ) 2 C0 3 lessens the solubility of the carbonates
of Ba , Sr , and Ca , while their solubility is increased by the presence of
NH 4 C1 . MgC0 3 is soluble in ammonium carbonate and in ammonium
chloride, so much so that in presence of an abundance of the latter it is
not at all precipitated by the former, i. e. (NH 4 ).,C0 3 does not precipitate a
solution of MgCl 2 as the NH 4 C1 formed holds the Mg in solution.

179. These metals may be all precipitated as phosphates in presence
of ammonium salts, but their further separation for identification or esti-
mation would be attended with difficulty (145 and //.).

180. The oxalates of Ba , Sr, and Mg are sparingly soluble in water,
calcium oxalate insoluble. Barium chromate is insoluble in water (27
and 186, 5^), strontium chromate sparingly soluble, and calcium and mag-
nesium chromates freely soluble.

181. In qualitative analysis, the group-separation of the fifth-group
metals is effected, after removal of the first four groups of bases, by

Online LibraryA. B. (Albert Benjamin) PrescottQualitative chemical analysis; a guide in qualitative work, with data for analytical operations and laboratory methods → online text (page 29 of 57)