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and this is normally present in all water that is exposed to the atmosphere,
can accomplish this oxidation of the ferrous ions to the ferric condition and it
also aids in the oxidation of the iron from the metallic to the ferrous state. In
this way iron exposed to moisture and oxygen oxidizes or rusts. The rusting
process is favored by the contact of the metal with a more noble metal, such as
platinum, copper or nickel; an electric couple is formed and the iron becomes
the positive pole, so that the hydrogen set free by the action of iron upon water
is deposited upon the more noble metal. The presence of a less noble metal,
such as zinc, tends to hinder the corrosion of iron the zinc corrodes instead of
the iron.

The presence of an acid is, therefore, not absolutely necessary to start the
corrosion of iron. An increase in the concentration of hydrogen ions, however,
will greatly hasten the solution of the metal. When carbonic acid is present,
ferrous bicarbonate is first formed and, when the ferrous iron is oxidized to the



144 REACTIONS OF THE METALS

4

ferric condition, the carbonic acid is set free again because ferric carbonate does
not exist. The acid again acts upon the metal and the rate of its corrosion is
greatly accelerated.

Certain substances tend to make iron passive, particularly strong nitric acid.
Passive iron does not dissolve in dilute nitric acid and does not corrode readily
(cf. aluminium, p. 126). On the other hand, certain substances can overcome
the passive condition and are said to activate the iron. Thus a solution of com-
mon salt is an activating agent.

Different varieties of iron and steel corrode with different degrees of readi-
ness. Cast iron is often protected by its casting skin. Impurities present in
steel often favor corrosion by causing electric couples to be established.

A. Ferrous Compounds

Ferrous compounds, which may be prepared by dissolving metallic
iron, ferrous oxide, ferrous hydroxide, ferrous carbonate, or ferrous
sulfide, etc., in acids, are usually greenish in the crystallized state,
but in the anhydrous condition they are white, yellow or bluish; in
concentrated solution they are green; in dilute solutions almost color-
less. Ferrous compounds exhibit a strong tendency to change over into
ferric salts; they are strong reducing agents.

REACTIONS IN THE WET WAY

1. Ammonia produces in neutral solutions an incomplete pre-
cipitation of white ferrous hydroxide :

FeCl 2 +2NH 3 +2H 2 O <= Fe(OH) 2 +2NH 4 Cl.

Ferrous salts in this respect are similar to those of magnesium
(cf. p. 94). In the presence of ammonium chloride the reaction takes
place in the direction from right to left; ammonia, therefore, causes
no precipitation with ferrous salts out of contact with the air, provided
sufficient ammonium chloride is present. On exposure to the air,
however, a turbidity is soon formed, green at first, then almost black,
and finally becoming brown. The small amount of ferrous hydroxide
contained in the solution is oxidized by the air, forming at first black
ferrous-ferric hydroxide and finally brown ferric hydroxide.

2. Potassium and Sodium Hydroxides produce, if air is excluded,
complete precipitation of white ferrous hydroxide,

Fe+20H - +Fe(OH) 2 ,

which is quickly oxidized by the air into ferric hydroxide.

3. Hydrogen Sulfide produces no precipitation in acid solutions
of ferrous salts; in dilute neutral solutions a small amount of black
ferrous sulfide is precipitated; but if the solution contains consid-



IRON 145

erable alkali acetate, hydrogen sulfide precipitates more of the iron as
ferrous sulfide (but not all of it), in spite of the fact that ferrous
sulfide is readily soluble in acetic acid. This interesting fact is an
instructive illustration of the law of chemical mass action.

The table on p. 22 states that 3.4X10~ 8 g. of ferrous sulfide dissolves
in a liter of water. This small quantity exists in solution almost entirely as
Fe ++ and S = ions. FeS <=* Fe ++ +S = . When acetic acid, which is a much
stronger acid than hydrogen sulfide (cf. p. 10), is added to the solution, equilib-
rium has to be established between its hydrogen ions and the dissolved sulfide
ions, 2H + +S = <=^ H 2 S, and, as a result of the formation of non-ionized hydrogen
sulfide, the solution no longer contains enough sulfur ions to reach the value of
the solubility product of FeS; to restore the equilibrium between FeS and its
ions, more of the solid must dissolve. If, moreover, the solution is boiled, the
hydrogen sulfide escapes as a gas as soon as it is formed. Consequently it is
impossible to arrive at a state of equilibrium until all of the ferrous sulfide has
dissolved. The solution is accomplished by means of hydrogen ions:

FeS+2H+ <= Fe ++ +H 2 S | .

On the other hand, if the ionization of the acetic acid is repressed by adding
an alkali acetate to the solution (cf. p. 46), and the concentration of the hydro-
gen sulfide is made as large as possible by keeping the solution saturated with
the gas, the reaction will take place in the reverse direction and some of the
iron will be precipitated as ferrous sulfide.

4. Ammonium Sulfide precipitates iron completely as black fer-
rous sulfide :

FeCl 2 + (NH 4 ) 28 = 2NH 4 C1 +FeS ,

which is readily soluble in acids with evolution of hydrogen sulfide.
In moist air it turns slightly brown, a part of the sulfur separates out,
and a basic ferric sulfate is formed.

5. Alkali Carbonates precipitate the white carbonate,

FeCl 2 +Na 2 CO 3 = 2NaCl+FeCO 3 ,
which in contact with the air becomes green, then brown :

4FeCO 3 +6H 2 0+0 2 = 4C0 2 +4Fe(OH) 3 ,

being converted into ferric hydroxide with loss of carbonic anhy-
dride.

Ferrous carbonate, like calcium carbonate (cf. p. 103), is soluble in
carbonic acid, forming ferrous bicarbonate:

FeC0 3 +H 2 CO 3 = FeH 2 (C0 3 ) 2 ,

a compound which is found in many natural waters, but which, like
the normal carbonate, is decomposed by atmospheric oxygen with
separation of ferric hydroxide:



146 REACTIONS OF THE METALS

Consequently a mineral water which contains ferrous bicarbonate,
if allowed to stand in contact with the air, will become turbid, owing
to the deposition of ferric hydroxide. To prevent this, the bottle
must be filled with water and tightly corked, so that no trace of air
can get in. Ferric hydroxide is insoluble in carbonic acid.

6. Potassium Cyanide precipitates yellowish-brown ferrous cyanide,

Fe ++ +2N(r-*Fe(CN) 2 ,

which is soluble in excess of the reagent, forming potassium ferro-
cyanide :

Fe(CN) 2 +4CN - Fe(CN) 6 == .

The complex ferrocyanide anion is in equilibrium, to be sure, with
simple ferrous cations, but the quantity of the latter present in the
aqueous solution of a ferrocyanide is so small that none of the above
reactions characteristic of ferrous ions can be obtained with it. Many
other similar complex cyanide anions are known; thus, the cyanides
of silver, nickel, iron (ferrous and ferric) , and cobalt all dissolve in potas-
sium cyanide, forming the following complex ions: [Ag(CN)2] ,
[Ni(CN) 4 ] = [Fe(CN) 6 ]", [Fe(CN) 6 ] = ' = [Co(CN) fl ] s . The acids are:
H[Ag(CN) 2 ], H 2 [Ni(CN) 4 ], H 3 [Fe(CN) 6 ], H 4 [Fe(CN) 6 ], H 3 [Co(CN) G ].
It is possible, as a matter of fact, to isolate the last three acids, though
the two former have never been prepared; they immediately break
down into metallic cyanide and hydrocyanic acid, just as carbonic acid
is decomposed into water and carbon dioxide.

With iron, therefore, there are two series of complex cyanogen
compounds, the ferrocyanides and the ferricyanides. The ferro-
cyanic derivatives contain the quadrivalent ferrocyanide anion and the
ferricyanides contain the trivalent ferricyanide anion.

Potassium ferrocyanide, K4[Fe(CN)e], is often called yellow prussiate
of potash, and potassium ferricyanide, Ks[Fe(CN)6], is called red prus-
siate of potash. The solubility of the alkali and alkaline-earth salts, and
the insolubility and color of the salts of the heavy metals (especially
with both ferric and ferrous iron), are very characteristic of ferro- and
ferricyanides.

7. Potassium Ferrocyanide, K 4 Fe(CN) 6 , produces in solutions
of ferrous salts, with complete exclusion of air, a white precipitate of
potassium ferrous ferrocyanide or of ferrous ferrocyanide, depending
upon whether one or two molecules of ferrous salt react with one mole-
cule of potassium ferrocyanide:

K 4 [Fe(CN) 6 ]+FeSO 4 = K 2 SO 4 +K 2 Fe[Fe(CN) 6 ]
= 2K 2 SO 4 +Fe 2 [Fe(CN)6]



IKON 147

Although both of the above salts are white, a light-blue color is almost
always obtained, because the precipitate is immediately oxidized some-
what by the air, forming the ferric salt of hydroferrocyanic acid (Prus-
sian blue) :



Prussian blue



8. Potassium Ferricyanide, Ks[Fe(CN)e], added to solutions of
ferrous salts produces a dark blue precipitate (Turnbull's blue) consist-
ing of ferrous ferricyanide mixed with potassium-ferric ferrocyanide :



and

K 3 [Fe(CN)6]+FeCl2=KFe[Fe(CN) 6 ] = =+2KCl.

Blue

In other words the ferricyanide acts both as a precipitant and as an oxidizing
agent * and a blue color results whenever iron is present in the cation in a state
of oxidation different from that of the iron present in the complex anion. The
ferricyanide ion is a strong oxidizing agent and in alkaline solution readily
oxidizes ferrous hydroxide to ferric hydroxide. Turnbull's blue is not very
soluble in acid solutions, but is decomposed by treatment with caustic alkali,
all of the complex anion being in the form of ferrocyanide:

Fe 3 ++ [Fe(CN) 6 ] 2 s +8KOH =2K 4 [Fe(CN) 6 ] +2Fe(OH) 3 +Fe(OH) 2 ,
K+Fe+ + +[Fe(CN) 6 l = == +3KOH = K 4 [Fe(CN) 6 ] +Fe(OH) 3 .

9. Potassium Thiocyanate gives no reaction with ferrou3 salts
(note difference from ferric salts).

As has been stated, ferrous salts are oxidized by the air to ferric salts; thus
ferrous sulfate is gradually changed into brown, basic ferric sulfate,

2FeS0 4 +0=Fe 2 0-(S0 4 ) 2 ,

which is insoluble in water. Consequently it often happens that ferrous sul-
fate will not dissolve in water to a clear solution, but gives a brown, turbid
solution, becoming clear on the addition of acid, the basic ferric salt being
changed to a soluble neutral salt:

Fe 2 (S0 4 ) 2 +H 2 S0 4 -> Fe 2 (S0 4 ) 3 +H 2 0.

Such a solution, which then contains ferric ions, reacts with potassium thio-
cyanate (cf. p. 150). To free the solution from ferric salt, it may be boiled
with metallic iron, with exclusion of ah*, whereby the ferric salt is changed into
ferrous salt:

Fe 2 (S0 4 ) 3 +Fe=3FeS0 4 .

By means of strong oxidizing agents, ferrous salts can be quickly and
completely changed into ferric salts, as was shown in the introduction (cf.
pp. 27-33).



Cf. ERICH MULLER, /. pr. Chem., 84 (1911), 353.



148 REACTIONS OF THE METALS

Detection of Ferrous Oxide in the Presence of Metallic Iron

Treat the mixture with a large excess of a neutral solution of mercuric
chloride and heat on the water-bath; the metallic iron goes into solution as
ferrous chloride:

2HgCl 2 +Fe =FeCl 2 +Hg 2 Cl 2 .

Filter off the residue and test the filtrate with potassium f erricyanide ; a
precipitate of Turnbull's blue shows that metallic iron was originally present.

Wash the residue with cold water, until all of the ferrous chloride has been
dissolved, and then treat it with dilute hydrochloric acid. If the solution
now gives a precipitate of Turnbull's blue with potassium ferricyanide, ferrous
oxide was present.

If hydrogen is given off, some metallic iron is still present; the experiment
must be repeated and the mixture given a longer treatment with HgCL solution.

B. Ferric Compounds

Ferric oxide, Fe20s, is reddish brown, becomes grayish black on
strong ignition, but on being pulverized appears red again.

The ferric salts are usually yellow or brown, but ferric am-
monium alum is pale violet. Ferric salts are yellowish brown in aque-
ous solution, and the solution reacts acid (hydrolysis). Dilution and
warming favor the hydrolysis, so that all strongly diluted ferric salts
deposit basic salts on being boiled:

Fe 2 (S0 4 ) 3 +H 2 <=Fe 2 (SO 4 ) 2 O+H 2 S0 4 .

With ferric salts of the weaker acids, often all of the iron is pre-
cipitated as a basic salt; thus the acetate, on being boiled in a dilute
solution, reacts as follows:

Fe(C2H 3 02)3+2H 2 ^Fe(OH) 2 (C 2 H 3 C>2) +2HC 2 H 3 2 .



By the addition of acid all basic salts may be changed back into
neutral salts.

REACTIONS OF FERRIC SALTS IN THE WET WAY

1. Ammonia precipitates brown, gelatinous ferric hydroxide:

Fe + + + +30H - Fe(OH) 3 .

The solubility product of ferric hydroxide is so small (cf. p. 22)
that it is precipitated completely even in the presence of ammonium
salts; it is readily soluble in acids. On ignition it loses water and is
changed to oxide, which is very difficultly soluble in dilute acids. It is
best brought into solution by long-continued heating below the boiling
point with concentrated hydrochloric acid.



IRON 149

2. Potassium and Sodium Hydroxide also precipitate ferric hy-
droxide.

3. Sodium Carbonate produces a brown precipitate of basic car-
bonate, which at the boiling temperature is completely decomposed
hydrolytically into hydroxide and carbon dioxide:

2FeCl3+3Na 2 CO 3 +3H 2 O = 2Fe(OH) 3 +6NaCl-f-3CO 2 T .

4. Zinc Oxide and Mercuric Oxide also precipitate the iron as
hydroxide :

2FeCl 3 +3ZnO+3H 2 O = 3ZnCl 2 +2Fe(OH) 3 .

This reaction is frequently used in quantitative analysis.

5. Sodium Phosphate precipitates yellowish-white ferric phosphate:

FeCl 3 +2Na 2 HPO4 = 3NaCl+NaH 2 PO4+FePO 4 .

Ferric phosphate is insoluble in acetic acid, but readily soluble
in mineral acids. The precipitation of iron with sodium hydrogen
phosphate is consequently only complete when a large excess of the
precipitant is employed, or when sodium acetate is added:



In this last case all the iron and all the phosphoric acid are pre-
cipitated. The reaction is often used to precipitate phosphoric acid
quantitatively. An excess of the disodium phosphate will also cause
complete precipitation of iron as phosphate, if the -phosphate solution
is previously exactly neutralized with ammonia:



and

Na 2 NH4PO4+FeCl 3 = 2NaCl+NH 4 Cl+FePO4.

If, however, an excess of sodium phosphate and ammonia is added
to the iron solution, the precipitation of iron is incomplete, because the
ferric phosphate dissolves in the excess of sodium phosphate, in the
presence of ammonia (or ammonium carbonate), with a brown color
and formation of a complex salt.

Ferric phosphate is transformed by ammonia into a brown basic
phosphate, and by potassium hydroxide almost completely into ferric
hydroxide and potassium phosphate; while by fusion with caustic
alkali or alkali carbonate it is completely decomposed.

If alkaline earth ions are present, an excess of ammonia completely
changes ferric phosphate to ferric hydroxide and alkaline curtli phos-
phate is precipitated.



150 REACTIONS OF THE METALS

6. Alkali Acetates produce in cold, neutral solutions a dark-brown
coloration, and on boiling the dilute solution all of the iron separates
as basic acetate:

FeCl 3 +3NaC 2 H 3 02 = 3NaCl+Fe(C 2 H 3 02) 3 (in the cold),

(on boiling).



The presence of organic hydroxy-acids (tartaric, malic, citric, etc.)
and of polyatomic alcohols (glycerol, erythritol, mannitol, sugars, etc.)
prevent all of the above-mentioned reactions, because complex salts
are formed in which the iron is present in the form of a complex anion
(cf. aluminium, p. 128).

7. Potassium Thiocyanate, KCNS, produces in solutions of ferric
salts a blood-red coloration:

Fe + + + +3CNS~ <=> Fe(CNS) 3 .

This action is reversible; the red color of the slightly ionized ferric thio-
cyanate being most intense when an excess of ferric salt, or of potassium thio-
cyanate is present.

If the solution is shaken with ether, the Fe(CNS) 3 goes into the ether. Ferric
thiocyanate combines readily with potassium thiocyanate, forming complex
potassium ferrithiocyanate:

Fe(CNS) 3 +3KCNS =K 3 [Fe(CNS)] 6 *

analogous to potassium ferricyanide, K 3 [Fe(CN) 6 ].

The complex salt is insoluble in ether, the Fe(CNS) 3 only being soluble
therein ; so that the red color is due to the formation of the ferric thiocyanate
and not to the complex salt.

This reaction is extremely sensitive, but cannot always be relied on. The
test cannot be made in the presence of strong oxidizing agents such as nitric
acid, as a red color is produced by the oxidation of the thiocyanate. The
oxidized compound is not very stable, however, and its color is not very deep.
If the solution contains considerable alkali acetate, the coloration cannot be
recognized. The presence of organic hydroxy-compounds (tartaric acid, etc.)
prevents the reaction in neutral solutions, but not in acid solutions. In the
presence of mercuric chloride the red color disappears entirely; the mercuric
chloride reacts with the ferric thiocyanate, forming a colorless, soluble
mercuric double salt, which is ionized even less than ferric thiocyanate:

2Fe(CNS) 3 +6HgCl 2 =2FeCl 3 +3[Hg(CNS) 2 .HgCl 2 ].

8. Potassium Ferrocyanide, K4Fe(CN)6, produces in neutral or
acid solutions of ferric salts an intense blue precipitation of Prussian
blue:



K 3 [Fe(CNS) 6 ]+4H 2 O. Cf. ROSENHEIM, Z. anorg. Chem., 27 (1901), 208.



IRON 151

Prussian blue, the ferric salt of ferrocyanic acid, is insoluble in water, but
soluble in oxalic acid and in an excess of potassium ferrocyanide; the solution
thus obtained is a deep blue and is used as blueing and as blue ink. The blue
solution obtained with a ferric salt and an excess of potassium ferrocyanide con-
tains colloidal KFe[Fe(CN) 6 ]-H 2 0, which can be salted out by the addition of a
considerable quantity of electrolyte such as alkali chloride, sulfate or nitrate.
Prussian blue is also soluble in concentrated hydrochloric acid, but is pre-
cipitated again on dilution. As the ferric salt of ferrocyanic acid it behaves
like other ferric salts to the hydroxides of the alkalies; ferric hydroxide and the
alkali salt of hydroferrocyanic acid being formed:

Fe 4 [Fe(CN) 6 ] 3 +120H- -> 4Fe(OH) 3 +3[Fe(CN) 4 f i .

9. Potassium Ferricyanide, Ka[Fe(CN)a], produces no precipita-
tion in solutions of ferric salts, only a brown coloration (differing
from ferrous salts) :



10. Ammonium Sulfide, added to a solution of a ferric salt, gives
a precipitate of ferric sulfide, Fe2Sa,

2Fe + ++ +3S=-+Fe 2 S 3 ,

which is soluble in cold, dilute hydrochloric acid, forming ferrous
chloride and sulfur:

Fe 2 S 3 +4H + -2Fe ++ -f 2H 2 S T +S.

The fact that Fe 2 S 3 is precipitated, and not FeS as commonly believed, was
proved by H. N. Stokes* who decomposed it out of contact with air by zinc-
ammonium oxide and obtained white ZnS and red Fe(OH) 3 . L. Gedelf has also
shown that hydrogen sulfide passed into a solution of ferric chloride made
alkaline with ammonia gives Fe 2 S 3 . If, however, the solution is acid, hydrogen
sulfide or ammonium sulfide reduces the iron before any precipitate is formed.

11. Hydrogen Sulfide in acid solutions reduces ferric salts to
ferrous salts, with separation of sulfur :

2Fe +++ +H 2 S -* 2Fe ++ +2H + +S.

Besides hydrogen sulfide, many other substances (nascent hydrogen,
stannous chloride, sulfurous acid, hydriodic acid, etc.) will reduce ferric
salts, as was shown on pp. 35, 36.

12. Ether when shaken with a solution of ferric chloride in 6 N
hydrochloric acid dissolves most of the ferric chloride. By separating



* J. am. Chem. Soc., 29 (1907), 304.

t Ueber Schwefeleisen, Karlsruhe, (1905).



152 REACTIONS OF THE METALS

the ether with the aid of a separately funnol, and repeating the opera-
tion, nearly all of the iron can be removed from the aqueous solution.
(Cf. p. 17.)

13. Cupferron, the ammonium salt of phenylnitrosohydroxylamine,
C 6 H 5 NO-NONH 4 , precipitates red (CeHsNO NO) 3 Fe, which is soluble
in ether, insoluble in acids, and converted into Fe(OH) 3 by treatment
with ammonia.

14. Sodium Thiosulfate, Na2S203, colors neutral ferric solu-
tions a violet red, but the color disappears quickly and the solution
then contains ferrous salt and sodium tetrathionate:

2Na 2 S 2 O 3 + 2FeCl 3 = 2NaCl + 2FeCl 2 +Na 2 S 4 06.

The composition of the violet-red substance which is first formed
is unknown; perhaps it is ferric ihiosulfate.

As we have seen, there exist a number of iron compounds which contain
the metal as a complex ion, so that it cannot be detected by the ordinary
reagents. The complex hydroxy-organic compounds, as well as the ferro-
and ferricyanide compounds, belong to this class of compounds.

If it is a question of proving the presence of iron in such a compound, a
different method should be used in the case of an organic hydroxy-compound
from that in the case of a ferro- or ferricyanide.

If organic substances are present, the iron is precipitated as sulfide by means
of ammonium sulfide; or the organic matter is first removed by ignition,
whereby metallic iron, oxides of iron and carbon are obtained.

In case we have a ferro- or ferricyanide, the iron cannot even be precipitated
by means of ammonium sulfide; the compound must be completely destroyed
before it will be possible to detect the presence of iron by any of the ordinary
methods.

This may be accomplished (a) by ignition, (6) by fusion with potassium
carbonate or sodium carbonate, or (c) by heating strongly with concentrated
sulfuric acid.

(a) Decomposition by Ignition. The ferrocyanides are decomposed (with
evolution of nitrogen) into potassium cyanide and carbide of iron:

K 4 [Fe(CN) 6 =4KCN+FeC 2 +N, T .

The ferricyanides also leave behind iron carbide and potassium cyanide,
but evolve cyanogen as well as nitrogen :

2K,[Fe(CN).]=6KCN+2FeC,+ (CN) a +2N a T .

Treat the residue from the ignition with water, whereby the potassium
cyanide goes into solution, leaving behind the iron carbide; filter and treat
the residue with hydrochloric acid. The iron goes into solution as ferrous
chloride, hydrocarbons are given off, and there remains some carbon.

The above decomposition can be imagined to take place as follows :



IRON 153

By heating potassium ferrocyanide, it is decomposed first into potassium
cyanide and ferrous cyanide, while the latter on further heating is changed
to iron carbide and nitrogen :

(a) K 4 [Fe(CN) 6 l=4KCN+Fe(CN) 2 ;
(0) Fe(CN) 2 =FeC 2 +N 2 T.

Potassium ferricyanide is decomposed into potassium cyanide and the
very unstable ferric cyanide, which splits off cyanogen and becomes ferrous
cyanide; the latter is decomposed, as before, into iron carbide and nitrogen:

(a) K 3 [Fe(CN) 6 ]=3KCN+Fe(CN) 3 ;
(/3) Fe(CN) 3 =Fe(CN) 2 +CN T ;
( 7 ) Fe(CN) 2 = FeC 2 +N 2 T.

(6) Decomposition by Means of Fusion with Potassium Carbonate. Mix the
substance with an equal amount of the carbonate and heat in a porcelain cru-
cible until a quiet fusion is obtained. By this means a mixture of potassium
cyanide and potassium cyanate (both soluble in water) is formed in the presence
of metallic iron:

K 4 [Fe(CN) 6 ]+K 2 C0 3 = 5KCN+KCNO+C0 2 T +Fe.

Extract the melt with water, filter and dissolve the iron in hydrochloric
acid.

(c) Decomposition by Heating with Concentrated Sulfuric Acid. By heating
with concentrated sulfuric acid all complex cyanogen compounds may be
decomposed. By this means the metal present is changed into sulfate, the
nitrogen of the cyanogen into ammonium sulfate, while the carbon of the
cyanogen escapes as carbon monoxide:

K 4 [Fe(CN) 6 ]4-6H,S04+6H 2 0=2K 2 S0 4 +FeS0 4 +3(NH 4 ) 2 S04+6CO T ,
2K 3 [Fe(CN) 6 ]+12H 2 S0 4 +12H 2 0=3K 2 S0 4 +Fe 2 (S0 4 )3+6(NH 4 ) 2 S0 4 +12CO T .

The treatment with concentrated sulfuric acid is best accomplished in a
procelain crucible placed in an inclined position over the flame, and the flame
directed against the upper part of the crucible. Continue heating until fumes
of sulfuric acid cease to come off. Treat the residue, which consists of an
alkali sulfate and anhydrous ferrous or ferric sulfate, with a little concen-
trated sulfuric acid, heat gently, and add water little by little. In this way the
sulfate is readily brought into solution.

REACTIONS IN THE DRY WAY

The borax (or sodium metaphosphate) bead, containing a small
amount of an iron salt, is yellow while hot and colorless when cold
after being heated in the oxidizing flame, and pale green after being
heated in the reducing flame. When strongly saturated, however, the
bead obtained with the oxidizing flame is brown while hot, yellow



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