of the salt =CH(OMe), thus disturbing all the conditions of
equilibrium and ‹leading to the transformation of a very much larger
part of the aldehyde into a compound containing the characteristic
unsaturated (bivalent) carbon›, than was originally present. The
aldehyde will thus become ‹more susceptible to oxidation› as a result
of the enormous increase in the concentration of the oxidizable
component. We may assume this to be either the salt, =CH(OMe), or
its negative ion, =CH(O^{−}), or both, or some analogous derivative.
Further, the two free valences of a bivalent carbon atom may be
considered to consist of a positive and a negative charge of
electricity, either actual or potential,[583] and ‹the oxidation›
will consist [p292] ‹primarily in the absorption of two positive
charges, from the oxidizing agent, to convert the negative charge on
the carbon atom›, say in ±CH(ONa), ‹into a positive charge›.[584]
If the oxidizing agent is alkaline silver nitrate solution, we may
formulate the successive actions as follows:
(NaO)HC± + 2 Ag^{+} → (NaO)HC^{2+} + 2 Ag ↓.
The two positive silver ions correspond to two negative ions, ‹e.g.›
hydroxide ions HO^{−}, which are set free by the discharge of the
silver ions, and which, in turn, will combine with the oxidized
carbon atom holding the two positive charges:
(NaO)HC^{2+} + 2 HO^{−} → (NaO)HC(OH)_{2} → (NaO)HC:O + H_{2}O.
The salt formed, HCO_{2}Na, is ‹sodium formate›, which is the first
isolated product of the oxidation of formaldehyde.
It would appear, from this point of view, that the ‹alkaline nature›
of the silver nitrate mixture is advantageous primarily because a
base is required by the formaldehyde, the reducing agent, to convert
it into some readily oxidizable form. And the proved efficiency
of the alkaline mixture (see above) makes it appear probable that
the advantage gained by this result ‹more than offsets the loss
in oxidizing power›, suffered by the silver nitrate following the
suppression of its real oxidizing component, the silver-ion, when,
in the presence of ammonia, the latter is converted largely into
the ion, Ag(NH_{3})_{2}^{+}. Ammonia, in turn, is employed in
the oxidizing mixture, essentially with the object of preventing
the precipitation of the silver-ion, as silver oxide, by the
hydroxide-ion of an alkaline mixture. These conclusions, as well
as, in particular, the ‹main conception› «that in the oxidation of
formaldehyde there is an actual transfer of electrical charges», may
be fully confirmed with the aid of the chemometer.[585]
EXP. A small beaker, containing a platinum electrode, which is
connected with the positive post of the voltmeter, is half filled
with a solution of silver and sodium nitrates. A similar small
beaker, containing a platinum electrode leading to the negative post
of the voltmeter, is charged with a solution of sodium nitrate (to
render the solution a good conductor) and with some formalin. The
solutions in the two beakers are connected by means of a salt-bridge
containing sodium nitrate. [p293]
Only a ‹very slight› current is produced under these conditions; the
potential between silver nitrate and formaldehyde is found to be
‹extremely› small. If, now, sodium hydroxide is added to the formalin
mixture, an ‹enormous increase› in potential is observed, proving,
unmistakably, that the addition of the alkali to the formalin
solution ‹enormously increases the concentration of the reacting,
oxidizable component›.[586]
When some ammonia is added to the silver nitrate mixture, we find,
as anticipated, that the ‹oxidizing power of the silver solution is
greatly reduced›, the silver-ion being converted into the complex
ion, Ag(NH_{3})_{2}^{+}; but the potential is still very much
‹greater› than the potential between silver nitrate and formalin
without any alkali—which shows that the advantage of using alkali
with the formaldehyde greatly outweighs the disadvantage of using
ammonia with the silver nitrate.
An electric current may also be readily obtained by combining
alkaline formaldehyde with other oxidizing agents—for instance with
an ‹oxygen› electrode (p. 279). We find (‹exp.›) that the oxidation
proceeds with remarkable ease under these conditions. Permanganate,
dichromate, etc., may be substituted for oxygen, with the same
general result.
«Summary.»—It is clear, then, that the oxidation of an ‹organic
substance› may readily be interpreted as consisting, ultimately, in
a transfer of electrical charges, of exactly the same nature, as is
found in the other oxidation reactions which we have considered,
[p294] and that the conditions for producing a maximum current,
as investigated with the aid of the chemometer, give us, again,
important guidance in following the course of the reactions. It is
needless to say that the oxidation of other organic compounds, such
as glucose, alcohol, etc., may be profitably studied from the same
point of view.
It also follows, from the conclusions reached, that, under proper
experimental conditions, electricity, in the form of a current,
must be capable of effecting the oxidation, or the reduction,
of organic as well as inorganic compounds (p. 252). Extended
investigations have, indeed, shown that electric currents belong to
the most important and efficient agents for this purpose, because
the oxidation, or the reduction, of the organic compound becomes
susceptible to the most exact control through the regulation of the
potentials used.[587]
«Tables and Summaries.»—In the first table the equilibrium
(solution-tension) constants of a number of metals and non-metallic
elements are given. The table is followed by brief explanations of
its meaning and a summary of some of its more important applications.
TABLE[A] OF EQUILIBRIUM (SOLUTION-TENSION) CONSTANTS (IN MOLAR
TERMS) AND OF POTENTIAL DIFFERENCES BETWEEN ELEMENTS AND THEIR IONS
IN UNIMOLAR AQUEOUS SOLUTIONS.
Element, Ion. E.P._{El.,Ion.} K_{Ion.}
K, K^{+}[B] (−2.92) 6E50
Na, Na^{+}[C] −2.44 2.5E42
Ba, Ba^{2+} (−2.54) 2.1E88
Sr, Sr^{2+} (−2.49) 4.0E86
Ca, Ca^{2+} (−2.28) 2.0E79
Mg, Mg^{2+} (−2.26) 4.1E78
Al, Al^{3+}[D] −0.999 ? 1.3E52
Mn, Mn^{2+} −0.798 5.7E27
Zn, Zn^{2+} −0.493 1.4E17
Cd, Cd^{2+} −0.143 9.5E4
Fe, Fe^{2+}[E] −0.122 ? 1.8E4
Co, Co^{2+}[F] +0.0138 ? 0.3314
Ni, Ni^{2+}[G] +0.108 ? 1.8E−4
Sn, Sn^{2+} <+0.085 <1.1E−3
Pb, Pb^{2+} +0.129 3.3E−5
H_{2}, H^{+}[H] +0.277 1.52E−5
Cu, Cu^{2+} +0.606 8.3E−22
As, As^{+++} <+0.570 <2.7E−30
Bi, Bi^{3+} <+0.668 <1.4E−35
Sb, Sb^{3+} <+0.743 <1.7E−39
Hg, Hg^{+} +1.027 1.38E−18
Ag, Ag^{+} +1.048 6E−19
Pt, Pt^{4+} <+1.140 5E−80
Au, Au^{3+} <+1.356 <1.8E−71
F_{3}, F^{−}[H] (+2.24) 9.0E88
Cl_{2}, Cl^{−}[H] +1.694 3.16E29
Br_{2}, Br^{−} +1.270 1.23E22
I_{2}, I^{−} +0.797 ? 7.26E13
O_{2}, HO^{−}[I] +0.698 1.36E12
TABLE NOTES:
A. The table is based on Wilsmore's compilation of solution-tension
potentials, ‹Z. phys. Chem.›, «36», 91 (1901).
B. Values in parentheses have been estimated by indirect
measurements.
C. G. N. Lewis, ‹J. Am. Chem. Soc.›, «32», 1467 (1910).
D. Values marked with? are uncertain.
E. Calculated from the data of Richards and Behr (‹Z. phys.
Chem.›, «58», 301 (1907)), who found the potential of iron against
0.5 molar FeSO_{4} to be −0.15 volt. The degree of ionization of
0.5 molar FeSO_{4} is taken as 22%. [Λ = 25.8 (Kohlrausch and
Holborn, ‹loc. cit.›, p. 152) and Λ_{∞} is taken as 117, as for
ZnSO_{4} (‹ibid.›, p. 200).] On account of the doubtful value for
the degree of ionization, the values in the table are marked?, but
the value found by Richards and Behr appears to be quite accurate.
F. Calculated from the data of Schildbach (‹Z. für Elektroch.›,
«16», 967 (1910)). The same uncertainty as to the degree of
ionization exists as that discussed in the previous footnote.
G. Calculated from the data of E. P. Schoch (‹Am. Chem. J.›, «41»,
208 (1909)). The same uncertainty as to the degree of ionization
exists as that discussed in footnote 5, p. 294.
H. The values for ‹gaseous› elements refer to the gases under one
atmosphere pressure.
I. The potential of oxygen at 18°, 760 mm., against an alkaline
solution in which [HO^{−}] = 1. K_{Ion} refers to the concentration
of HO^{−}, with which oxygen under atmospheric pressure would be
directly in equilibrium, at 18°.
The potential differences, given in the table, are based on the
assumption that the ‹absolute zero› of potential is at such a
point, that the so-called standard normal calomel electrode has
a value of +0.56 volt relative to this zero (‹cf.› Ostwald, ‹Z.
phys. Chem.›, «36», 97 (1901)). The exact determination of this
value is a very difficult matter. Recently Palmaer (‹ibid.›, «59»,
129 (1907)), located the absolute zero at a point 0.04 volt more
positive than the above, making the absolute potential of the normal
calomel electrode, approximately, +0.52 volt. To refer potentials,
given in this book, to this new zero, one would subtract 0.04
volt from all positive potentials and add 0.04 to the numbers
representing negative potentials (‹e.g.› E. P._{Zn,Zn^{2+}} would
become −0.569 in place of −0.529 volt). Since the ‹equilibrium
(solution-tension) constants› are calculated from the potential
differences referred to the absolute zero (p. 259), any change
in the zero involves corresponding changes in the values of the
equilibrium constants, as calculated for this book. However, it
should be noted that all potential differences would be corrected
by the same constant quantity (0.04 volt for Palmaer's zero): ‹all
the equilibrium constants for univalent metallic ions would be
increased proportionally to a constant factor› ‹c› (‹c› is very
nearly equal to 5, for Palmaer's zero), the equilibrium constants
for bivalent metallic ions would be increased proportionally to
‹c›^2, etc. ‹The equilibrium ratio› for two metals and their
ions ‹would in no wise be changed› by these alterations: ‹e.g.›
for the equilibrium between zinc and copper and [p296] their
ions (p. 267), K_{equil.} = K_{Zn^{2+}} / K_{Cu^{2+}}; the
factor ‹c›^2 would be introduced into both terms of the ratio
and would not affect the value of the latter. For the condition
of equilibrium between silver and copper and their ions (p. 267)
K_{equil.} = K_{Ag^{+}}^2 / K_{Cu^{2+}}, and since (‹c›)^2 = ‹c›^2,
this equilibrium ratio would also not be affected. For elements,
which produce negative ions, the corresponding correction factors
would be 1 / ‹c›, 1 / ‹c›^2, etc., and the equilibrium relations
between two such elements and their ions likewise would remain
unchanged. Since these ‹equilibrium relations› are the ‹significant›
ones in this work, and since our conclusions have been based on
them, it is clear that a change in the absolute zero would not
affect the conclusions reached.
On account of the uncertainty attaching to the determination of
the absolute zero of potential, it is preferred, in practice,
to report the experimentally determined potentials as measured
against a constant, well-defined electrode (such as the calomel
electrode or a hydrogen electrode) and thus to eliminate the
variation, which a change in the determination of the zero potential
would make necessary. However, for an elementary discussion of
oxidation-reduction reactions, from the same viewpoint as is used
in considering all other reversible chemical actions, the idea of
the absolute potential has certain advantages, making a uniform
treatment possible.
1. ‹Meaning of K_{Ion}.› Under K_{Ion} is given, for each element,
the ‹concentration of its ion›, with which the element would be
directly in equilibrium at the ordinary temperature (see p. 258).
The constants for ‹gaseous› elements represent the constants of the
gases under atmospheric pressure.
2. ‹The Condition for Equilibrium between Two Elements and Their
Ions.› The condition of equilibrium in a system of two elements
and their ions may be found with the aid of the constants K_{Ion},
as follows: For Zn ↓ + Cu^{2+} ⥂ Zn^{2+} + Cu ↓ we have for the
condition of equilibrium (see p. 267)
[Zn^{2+}] / [Cu^{2+}] = K and K = K_{Zn^{2+}} / K_{Cu^{2+}} =
1.4E17 / 8.3E−22 = 1.7E38.
Zinc-ion must be present in enormous excess in the condition of
equilibrium and zinc will precipitate copper from solutions of
cupric salts until this relation is established. The ‹suppression
of the cupric-ion›—by precipitation in the form of insoluble salts
or by conversion into very stable complex ions—makes [Cu^{2+}]
exceedingly small and makes it increasingly difficult for zinc to
precipitate copper, and, under certain conditions, the ‹ordinary›
course of the action may be ‹reversed› (p. 268).
For Cu ↓ + 2 Ag^{+} ⥂ Cu^{2+} + 2 Ag ↓, we have (p. 267)
[Cu^{2+}] / [Ag^{+}]^2 = K and
K = K_{Cu^{2+}} / K_{Ag^{+}}^2 = 8.3E−22 / (6E−19)^2 = 2.3E15.
3. ‹Potential Differences Calculated with the Aid of K_{Ion}.› For
‹metallic elements›, which send out ‹positive ions›, in contact with
an aqueous solution containing the ion in concentration [C], the
potential difference is (see p. 261)
ε_{‹Me, Ion›^{‹v›}} = (0.0575 / ‹v›) ‹log›([C] / K_{Ion}) ‹volts›,
[p297]
and for elements which form negative ions (see footnote 1, p.
261),
ε_{‹Elem., Ion›^{‹v›}} = (−0.0575 / ‹v›) ‹log›([C] / K_{Ion}) ‹volts›.
In these equations ‹v› represents the valence of the ion. It is
clear that for the condition of equilibrium, in which [C] = K_{Ion},
the potential is 0. Further, for the potential difference between
copper and a cupric salt solution in which [Cu^{2+}] = 1, we would
have
ε_{Cu, Cu^{2+}} = (0.0575 / 2) log(1 / 8.3E−22) =
21.08 × 0.0575 / 2 = +0.606 volts.
4. ‹Meaning of E.P._{Element, Ion}.› Under E.P._{Element, Ion}
the table gives the ‹potential difference in volts›, calculated
for the ‹element› named and ‹an aqueous solution of its ion in
unit concentration› (one gram-ion per liter). For instance, for
zinc and [Zn^{2+}] = 1 (65.4 grams zinc-ion per liter), we have a
potential E.P._{Zn, Zn^{2+}} = −0.493. The signs used, in accordance
with the convention adopted (p. 261), indicate the ‹character of
the charge› on the ‹element electrode› (which is named ‹first› in
the subscript to E.P.). For instance, zinc in a solution in which
[Zn^{2+}] = 1 would acquire a ‹negative› charge (p. 266), the
potential difference E.P._{Zn, Zn^{2+}} being −0.493 according to
the table; silver, immersed in a solution in which [Ag^{+}] = 1,
would acquire a ‹positive charge›, the potential difference
E.P._{Ag, Ag^{+}} = +1.048.
The potentials given for the ‹gaseous› elements represent the
potentials of the gases under 760 mm. pressure.
5. ‹Potential Differences Calculated with the Aid of E.P._{Element,
Ion}.› The potential corresponding to any concentration [C] of a
‹metal ion› may be found from the equation[1]
ε_{El., Ion^{‹v›}} = E.P._{El., Ion} + (0.0575 / ‹v›) log[C] ‹volts›,
and the potential for any concentration [C] of the ions of elements
forming ‹negative ions› is found[588] according to
ε_{El., Ion^{‹v›}} = E.P._{El., Ion} − (0.0575 / ‹v›) log[C] ‹volts›.
6. ‹The Condition for Equilibrium between Two Metals and Their Ions,
Calculated with the Aid of E.P._{Element, Ion}.› The condition
for equilibrium in a system of two metals and their ions is
determined by the fact that the potential of the system must be 0
when equilibrium is established. We have, for instance for the two
metals zinc and copper and their ions, Zn^{2+} and Cu^{2+}, for
Zn ↓ + Cu^{2+} ⇄ Cu ↓ + Zn^{2+} the condition for equilibrium that
ε_{Cu, Cu^{2+}} − ε_{Zn, Zn^{2+}} = 0. According to the equation
given in § 5, we have, then, for the condition of equilibrium,
E.P._{Cu, Cu^{2+}} + (0.0575 / 2) log[Cu^{2+}] − E.P._{Zn, Zn^{2+}} +
(0.0575 / 2) log[Zn^{2+}] = 0.
[p298]
Then
(0.0575 / 2) log([Zn^{2+}] / [Cu^{2+}]) =
E.P._{Cu,Cu^{2+}} − E.P._{Zn,Zn^{2+}} = +0.606 − (−0.493) = 1.099.
From the last relation we find log([Zn^{2+}] / [Cu^{2+}]) = 38.2261,
and therefore, for the condition of equilibrium, [Zn^{2+}] /
[Cu^{2+}] = 1.7E38.
7. ‹Equilibrium Constants for Elements with Variable Concentration.›
The concentration of a pure metal at a given temperature may be
considered a constant, except in the case of extremely thin films
of the metal (p. 258). The concentration of hydrogen, and of the
non-metallic elements given in the table, is variable, and K_{Ion}
has a definite value only when the concentration of the element
is defined (see the preceding table, footnotes 3, 4, p. 295). For
certain estimations the ‹equilibrium constants›, which show the
relation between the two variables, namely the concentration of
the element and that of its ion, are very helpful (see pp. 274 and
275). In the following table some of the more important equilibrium
constants of this nature are given.
TABLE OF EQUILIBRIUM CONSTANTS.
Element. K_{equil.}.
Hydrogen: [H^{+}]^2 : [H_{2}] 5.6E−9
Oxygen: [HO^{−}]^4 : [O_{2}] 8.2E49
Chlorine: [Cl^{−}]^2 : [Cl_{2}] 2E60
Iodine[A]: [I^{−}]^2 : [I_{2}] 5.6E29
TABLE NOTE:
A. The value of the constant, as given, is only an approximate
estimation (p. 273).
The significance of the constants is indicated by the ratios given
in the table. The relation of these constants to those given in
the first table may be seen from the following illustration. For
hydrogen we have H_{2} ⇄ 2 H^{+}. The first table tells us that
hydrogen, at 18° under atmospheric pressure, is in equilibrium
with its ion when the concentration of hydrogen-ion is 1.52E−5
(under K_{Ion}). Now, a mole of hydrogen at 18° occupies
22.4 × 291 / 273 = 23.9 liters under atmospheric pressure, and
its concentration (per liter) is therefore 1 / 23.9 mole. Then
equilibrium exists, when [H_{2}] = 1 / 23.9 and [H^{+}] = 1.52E−5
and K_{equil.} = [H^{+}]^2 : [H_{2}] = (1.52E−5)^2 × 23.9 = 5.6E−9.
FOOTNOTES:
[553] These constants are calculated from data given in
Wilsmore's tables (‹loc. cit.›) on the solution tension of
hydrogen. Hydrogen, at 18° under one atmosphere pressure,
produces a potential of ε_{H_{2}, H^{+}} = +0.277 (see p.
261, in regard to the sign) against a solution containing
hydrogen-ion in a concentration [H^{+}] = 1 (see the table at
the end of this chapter). Now, there must be some concentration
of hydrogen-ion, which we will call [C], with which hydrogen
at 18° and 760 mm. pressure is directly in equilibrium,
with the potential 0. For any concentration of hydrogen-ion
[H^{+}], ‹other than› [C], a potential is produced according to
ε_{H_{2}, H^{+}} = 0.0575 log([H^{+}] / [C]). If we insert into
this equation the values [H^{+}] = 1 and the potential ε = +0.277,
and if we solve the equation for [C], we find [C] = 1.52E−5.
That is the concentration of H^{+}, with which hydrogen of
one atmosphere pressure at 18° is directly in equilibrium.
Since under these conditions of temperature and pressure
[H_{2}] = 1 / 23.9 mole, we have for the condition of ‹equilibrium›
[H^{+}]^2 / [H_{2}] = K : (1.52E−5)^2 : (1 / 23.9) = 5.55E−9 = K.
[554] Experimentally the relations for an "oxygen electrode" are
much more complicated than for a hydrogen electrode, as a result,
apparently, of the oxidation of the metal (‹e.g.› platinum), with
the aid of which the electrode is prepared. For a critical review
and summary of the more recent results on this point, ‹vide›
Schoch, ‹J. phys. Chem.›, «14», 665 (1910). For the purposes of
this book it will be sufficient to limit our discussion to the
behavior of an ideal oxygen electrode.
[555] The bivalent oxygen ions, O^{2−}, combine with hydrogen ions
(formed, for instance, by the ionization of water) and form the
more stable hydroxide ions (p. 246): O^{2−} + H^{+} + HO^{−} ⇄
2 HO^{−}, or simply, O^{2−} + H^{+} ⇄ HO^{−}. Then, [O^{2−}] ×
[H^{+}] / [HO^{−}] = k and [O^{2−}] = k × [HO^{−}] / [H^{+}].
But since we have [H^{+}] × [HO^{−}] = K_{HOH} for the ionization
of water (p. 176), we also have:
[H^{+}] = K_{HOH} / [HO^{−}] and
[O^{2−}] = (k / K_{HOH}) × [HO^{−}]^2.
By substituting this value for the concentration [O^{2−}] of the
oxide-ion in equation (1), equation (2) is obtained. The constant
K_{2} includes then the constants k and K_{HOH}.
[556] The constants are calculated from the estimated potential
of the oxygen-hydrogen cell, +1.231 volt, at 18°. (‹Vide›
G. N. Lewis, ‹Z. phys. Chem.›, «55», 465 (1906); Nernst and
Wartenberg, ‹ibid.›, «56», 534 (1906); Brönsted, ‹ibid.›,
«65», 91 (1908); and a summary and discussion by Schoch, ‹loc.
cit.›) At 18° oxygen, under one atmosphere pressure, gives an
estimated potential ε_{O_{2}, HO^{−}} = +1.508 against an ‹acid›
solution, in which the concentration of the ‹hydrogen-ion›
[H^{+}] = 1 (see the table at the end of this chapter). Since
at 18° [H^{+}] × [HO^{−}] = 0.81E−14, the value for [HO^{−}] in
this acid solution is 0.81E−14. Now, for oxygen, at 18° and 760
mm., there must be some concentration of hydroxide-ion, which we
will call [C], at which the tendency of the oxygen to ionize is
exactly balanced by the tendency of the hydroxide-ion to form
oxygen—at this point the potential is 0. For any concentration
[HO^{−}] of the hydroxide-ion, ‹other than› [C], a potential
will exist ε_{O_{2}, HO^{−}} = 0.0575 log([C] / [HO^{−}]). Since
for [HO^{−}] = 0.81E−14, we have a potential ε_{O_{2}, HO^{−}} =
+1.508, these values can be introduced into the equation and
the latter solved for [C]. We find thus [C] = 1.36E12, and
‹oxygen, at 18° and› 760 mm. ‹pressure, would be directly in
equilibrium with a solution in which› [HO^{−}] = 1.36E12. At
18° and 760 mm. pressure a liter of oxygen contains 1 / 23.9
mole, and thus we have for the condition of equilibrium
[HO^{−}]^4 : [O_{2}] = K : (1.36E12)^4 : (1 / 23.9) = 8.2E49 = K.
[557] The most convenient form of electrode for this purpose
consists (see Fig. 14) of a cylinder (about one inch long) of
platinum gauze, which is fused to a glass tube and connected with
a wire leading through the tube to some mercury, held in a small
branch tube, fused into the main tube near its upper end. The gas
is easily conducted to the platinum gauze electrode through such
a tube. The cylinder of platinum gauze may be made by joining the
ends of rolled gauze with pieces of molten glass. It is coated with
platinum black.
[558] ‹Cf.› footnote 1, p. 278.
[559] ‹Cf.› footnote 1, p. 279.
[560] See Ostwald, ‹Z. phys. Chem.›, «11», 521 (1893), Arrhenius,
‹ibid.›, «11», 805 (1893), and Nernst, ‹ibid.›, «14», 155 (1893),
for a detailed discussion of oxygen-hydrogen gas cells. For more
recent work, ‹vide› G. N. Lewis, ‹J. Am. Chem. Soc.›, «28», 158
(1905), where references to other recent investigations are given.
[561] ‹Cf.› footnote 1, p. 261, and pp. 277–279.
[562] See pp. 42, 252, for the expression of the changes as
‹transfers› of ‹electrons›.
[563] Fredenhagen, ‹Z. anorg. Chem.›, «29», 424 (1902), has
brought interesting experimental evidence of the charging of
an electrode with ‹gaseous oxygen›, when ferric-ion is the
oxidizing agent in aqueous solutions. Whether the oxygen,
which is liberated by the action of the ferric-ion on water,
4 Fe^{3+} + 4 HO^{−} ⇄ 4 Fe^{2+} + O_{2} + 2 H_{2}O, is always ‹the
intermediate product and the direct oxidizing agent› in aqueous
solution, can hardly be considered decided by the experiment—it
may well be the product of a ‹parallel action›, which must take
place to a certain extent, according to the laws of equilibrium,
in a system containing both Fe^{3+} and HO^{−} ions. The result
hardly proves that oxygen must be the intermediate product in
the ‹main action›, when ferric ions act as the oxidizing agent.
We may consider, for instance, a solution containing an iodide
and a ferric salt: iodide ions have a far smaller affinity for
their negative charges (electrons) than hydroxide ions have, and,
consequently, will transfer their negative charges (electrons)
more readily to the ferric ions than the hydroxide ions would. The
action, if oxygen were ‹first› liberated, would lead to the same
ultimate result, but the observation made by Fredenhagen would not
prove that the main action would not nevertheless go by the shorter
direct path rather than through an intermediate formation of oxygen.
[564] According to the theory, that arsenic acid is an oxidizing
agent because it gives up oxygen of a definite pressure, ‹this
pressure would be the more effective, the more completely the
opposing hydroxide-ion is suppressed› by added acid (p. 280; see
also p. 272, on the action of ferro-ion on silver-ion in the
presence and in the absence of fluorides).
[565] Only the simplest form of basic ionization of arsenic acid
is considered. Intermediate ionization into As(OH)_{4}^{−},
As(OH)_{3}^{2−}, etc. (see p. 249), is, of course, to be assumed in
any complete investigation of the subject.
[566] Arsenious acid As(OH)_{3}, or HAsO_{2}, as well as its anions
AsO_{3}^{3−} and AsO_{2}^{−}, may have their own characteristic
tendencies to assume positive charges and be oxidized to arsenic
acid and its derivatives (see footnote 1, p. 270). In ‹alkaline›
solutions these tendencies, and the potentials corresponding to
them, might well be more important factors in determining the
course of an action, than the tendency of As^{3+} to form As^{5+}.
The discussion in the text, which deals primarily with acid
solutions, does not exclude such relations.
[567] A millivoltmeter is used for this experiment.
[568] See footnote 1, p. 284, in regard to the interpretation of
this experiment on the basis of the theory of liberation of oxygen
by arsenic acid.
[569] ‹Cf.› Smith's ‹General Inorganic Chemistry›, p. 712.
[570] The essential feature of this point of view was first
published by Abegg, ‹Z. anorg. Chem.›, «39», 330 (1904), and ‹Z.
phys. Chem.›, «43», 385 (1903); ‹vide› also Stieglitz, ‹Am. Chem.
J.›, «39», 51 (footnote) (1908), and ‹Qualitative Analysis Notes›,
University of Chicago (1905). Abegg's view has recently received
support from J. J. Thomson in his ‹Corpuscular Theory of Matter›,
p. 118.
[571] H^{+} does not change its valence (charge) in the action,
and yet it appears as an essential component in both styles of the
current equations for the oxidation-reduction reaction.
[572] A somewhat similar development of these relations, for
arsenic acid, has been found in Abegg's ‹Anorg. Chem.›, III, 3, p.
552 (1907).
[573] Luther and Michie, ‹Z. für Elektroch.›, «14», 826 (1908).
[574] The calculation is based on the results obtained by Luther
and Michie.
[575] The proportion of UO_{2}^{2+} converted into U^{6+} would
be so minute, that in an experimental determination of the
concentration [UO_{2}^{2+}], the concentration [U^{6+}] would no
doubt be a negligible quantity.
[576] Peters, ‹Z. phys. Chem.›, «26», 193 (1898).
[577] In other instances, the action of the hydrogen-ion in
facilitating and accelerating chemical actions (often called its
"catalytic effect") has been explained, in a similar fashion, as
being based on ‹salt formation, followed by the ionization of the
salts formed›, the ‹active› components being the ions (‹vide›,
for instance, Bredig, ‹Z. für Elektroch.›, «9», 118 and «10»,
586 (1904) and Stieglitz, ‹Report of the Congress of Arts and
Sciences›, St. Louis, «4», 276 (1904) and ‹Am. Chem. J.›, «39»,
29, 415 (1908) and later articles). In many of these cases, the
ion concentrations of the reacting components have not yet been
accessible to direct measurements, but the viewpoint has been
sustained by quantitative studies of analogous reactions, which
were selected for study, because the factors involved could be
measured (‹cf.› Stieglitz, ‹loc. cit.›).
[578] Usher and Priestley, ‹Proc. Roy. Soc.›, B, «77», 369 (1905);
«78», 318 (1908).
[579] Tollens, ‹Ber. d. chem. Ges.›, «15», 1635 (1882).
[580] [Ag^{+}] × [NH_{3}]^2 / [Ag(NH_{3})_{2}^{+}] = 1 / 10^7.
[581] Methylene CH_{2}, itself, has never been isolated, but
derivatives of it are known, such as the cyanides, C(NH), C(NK)
(see pp. 66, 237).
[582] Potassium cyanide, =C(NK), is a powerful reducing agent (see
p. 89).
[583] Pp. 65, 66.
[584] If a negative charge is an electron, a positive charge the
absence of an electron in an atom, the bivalent carbon atom ‹loses
two electrons›, when it is oxidized.
[585] See Stieglitz, ‹Science›, «27», 774 (1908).
[586] The oxidation occurs essentially in the same manner as
described (p, 292) for the action of formaldehyde on ammoniacal
silver solution, when they are brought together in a single vessel.
In the present case, where the action is used to produce an
electric current, there is a ‹migration› of ‹negative ions› into
the formalin solution through the salt-bridge (p. 254). For every
two silver ions discharged on the electrode in the silver nitrate
solution, two hydroxide ions are liberated in the formaldehyde
solution, as a result of this migration, and they combine with the
oxidized carbon atom. The oxidation may be expressed, then, simply
as follows:
(NaO)HC± + 2 ⊕ + 2 HO^{−} → (NaO)HC^{2+} + 2 HO^{−} →
(NaO)HC:O + H_{2}O.
By comparison with the equations given on p. 292, it is evident,
that the only difference lies in the fact that the ‹positive
charges›, in the present case, are carried to the formaldehyde salt
through ‹metal wires› and a ‹metal electrode›, while previously
they were discharged ‹directly by silver ions on the formaldehyde
salt›.
[587] ‹Cf.› Nernst, ‹Theoretical Chemistry› (1904), p. 739, and the
applications mentioned there.
[588] Le Blanc, ‹Elektrochemie,› p. 215; ‹v› is the valence of the
ion.
[p299]
CHAPTER XVI
«SYSTEMATIC ANALYSIS FOR ACID IONS»
The systematic analysis for acid ions is made on a plan differing in
an important particular from the systematic analysis for metal ions.
The latter, as has been seen, are divided into groups, which, by
precipitation or solution of characteristic salts, are successively
separated from subsequent groups, before the isolated groups are
analyzed. That is, in general, a group of metal ions is examined
in the absence of the ions of all other groups. Acid ions are also
divided into groups, but, as a rule, the groups are not separated
from each other for analysis. The reason for this difference in
procedure is found, chiefly, in the fact that foreign acid ions
interfere[589] to a smaller degree with the specific tests for the
ions of a group, than is the case in the analysis for metal ions.
«Grouping of Acid Ions.»—While there is general agreement, among
the most important systems of analysis, in the grouping of metal
ions, there is notable variation in the way in which acids are
grouped by different authors. We shall confine ourselves here to the
consideration of two different modes of general procedure and not
discuss differences in minor details.
If acids were present only in the form of the free acids or of their
alkali salts, the division into groups could naturally and profitably
be made to include groups, which are identified by reactions carried
out in ‹neutral› or ‹alkaline› solution, as well as by such as
are made in ‹acid› solutions. Now, cations other than the alkali
ions are liable to interfere with tests designed for alkaline or
neutral solutions. For instance, a group of acid ions, in which
the phosphate-ion is included, is characterized by the fact that
the acids form barium salts, which are soluble in acid but not in
neutral or alkaline solutions. The absence or presence of such a
group may be recognized, if no cations other than the alkali metal
ions are present, by the addition of barium chloride to the solution
and by careful neutralization of any free acid, by ammonium [p300]
hydroxide. Barium phosphate and the barium salts of the other acids
of the group will be precipitated, if their ions are present. It is
clear, however, that a number of metal ions must interfere with the
test. For instance, a solution of aluminium nitrate or of ferric
chloride, treated with barium chloride, and with ammonium hydroxide
to neutralize the acid present in the solution as a result of the
hydrolysis of the salt, would give a precipitate of aluminium
hydroxide or of ferric hydroxide, and ‹not of barium salts›. The
formation of a precipitate under these conditions evidently will not
constitute any basis whatever for reaching a conclusion as to the
presence or absence of acid ions, such as can form precipitates of
barium salts under the same circumstances.
«Systematic Analysis for Acid Ions, Based on the Removal of Metal
Ions other than the Alkali Metal Ions.»—On account of this kind of
interference, by cations other than those of the alkali group, with
a number of group and specific tests for anions, that may be made in
neutral or alkaline solutions, provision is made, in most systems of
analysis, for the removal of such ions by proper treatment of the
substance under examination with sodium carbonate. Interfering metal
ions are thereby converted into carbonates or hydroxides which are
insoluble in water, while the acids form sodium salts which pass into
solution in water. Occasionally, recourse is also taken to hydrogen
sulphide to remove ions of the arsenic and copper groups.
The treatment with sodium carbonate, while advantageous in certain
cases, is ‹not uniformly successful›[590]: it is also frequently
complicated by the presence of amphoteric bases or of organic
substances, and frequently demands treatments and tests beyond those
made with the solution thus prepared. Furthermore, if a group is
found to be present, say a group of acid ions forming barium or
calcium salts insoluble in water, but soluble in acids,[591] all
the acids of the group, which have not been found to be present or
absent in the analysis for metal ions,[592] must be specifically
[p301] tested for, ‹although they may all be absent› and the only
representative of the group present may be an ion previously found in
the analysis for metal ions.[593] Then, too, when the group is found
to be absent by the group test, more sensitive tests for some of the
acid ions of the group must still be made to insure their complete
absence.[594]
While much may be said in favor of the systematic analysis for
acid ions, based on the preparation of a solution containing only
the alkali salts of the anions, and ‹while one should be familiar
with the plan and be able to have recourse to it at will›, yet the
drawbacks mentioned suggest another basis for the analysis.
«Systematic Analysis for Acid Ions in Acid Solution.»—If the
systematic analysis for acid ions is carried out entirely in ‹acid
solutions›, interference of cations with tests for anions is rarely
met with and in those rare cases may be easily provided against.
Such a method of systematic analysis in acid solution is frequently
more direct and more convenient than a method based on the removal
of cations other than the alkali metal ions. Almost all of the most
characteristic tests for anions, as it is, are carried out in acid
solutions, and only a very few good tests, which must be made in
neutral or alkaline solution, are sacrificed by placing the emphasis
on those carried out in acid solutions. The method, on the whole,
has proved a time-saving and convenient one, without loss in the
trustworthiness of the results.
«Desirability of Experience with Both Methods.»—It is desirable to
have experience with both methods, and to learn by such experience,
when to have recourse to the one or the other method. Suggestions as
to the choice of method are given in the Laboratory Manual, p. 119.
«The Groups of Acid Ions.»—The arrangement of the acid ions into
groups, for analysis in acid solution, does not differ in any
essential respect from the arrangement based on the use of solutions
of sodium salts of the acid ions—only one group test, which must
be made in neutral solutions, is omitted when the acid solution is
used, and the individual members of this group are tested [p302]
for, specifically. As that is also very frequently necessary
when solutions of sodium salts are used, no notable sacrifice in
convenience is made.
Since the grouping for both methods of analysis may be made the same,
the following grouping of acid ions has been adopted. Only the group
characteristics are given; the members of the groups are described in
detail in the Laboratory Manual (Part III).
I. ‹Ions of Amphoteric Acids and of Related Acids.› This group
includes those acids whose amphoteric character, or whose ready
reduction by hydrogen or ammonium sulphide, leads to their being
found, or indicated, in the systematic analysis for metal ions.
II. ‹The Carbonate Group.› This group includes those acids whose
physical properties (insolubility), or the physical properties of
decomposition products of which (carbon dioxide is a decomposition
product of carbonic acid), usually lead to their discovery in the
course of the preparation of solutions or of the analysis for cations.
III. ‹The Sulphate Group.› The ‹barium salts› of this group of acid
ions are ‹insoluble in acid solution›. Barium nitrate, added to a
solution acidified with nitric acid, is the group reagent.
IV. ‹The Chloride Group.› The anions of this group form ‹silver
salts›, which are ‹insoluble in nitric acid›. Silver nitrate, added
to a solution acidified with nitric acid, is the group reagent.
V. ‹The Phosphate Group.› A test for this group, as a whole, can be
made only if cations other than the alkali metal ions are absent:
the ‹barium salts› of the acid ions of the group are ‹insoluble
in neutral›, but ‹soluble› in strongly ‹acid, solutions›. Barium
nitrate, used with a neutral solution, is the group reagent in the
absence of metal ions other than the alkali metal ions. In the
presence of other cations, the three members of the group, which
are not found in some other group,[595] namely: phosphate, borate
and fluoride ions, are tested for specifically, and the group test
omitted. Phosphate-ion is tested for in nitric acid solution, in the
same solution as is used for the tests for groups III and IV. [p303]
VI. ‹The Nitrate Group.› The salts of the acids of this group are
readily soluble in water and specific tests for the acid ions are
made; there is no group test.
VII. ‹The Group of Organic Acids.› This group need only be considered
when a test for organic matter reveals its presence.
«Applications of Physico-Chemical Principles and Theories.»—The
physico-chemical principles and theories, which have been developed
in the previous chapters, naturally apply also to the reactions by
which acid ions are identified. In many cases such characteristic
reactions are identical with reactions studied in connection with
the metal ions. For instance, the precipitation of silver chloride,
used to identify the silver-ion (reagent, chloride-ion), may be used,
with certain precautions, to identify chloride-ion as well (reagent,
silver-ion).
In the following, only a few typical and interesting applications
of the principles and theories to acid ions will be given; numerous
other applications will suggest themselves in connection with the
laboratory work on the acids.
«Fractional Precipitation of Salts with a Common Ion.» For saturated
solutions of silver chloride, bromide and iodide, we have, according
to the principle of the solubility-product[596]:
[Ag^{+}]_{1} × [Cl^{−}]_{1} = K_{AgCl} = 1E−10;
[Ag^{+}]_{2} × [Br^{−}]_{2} = K_{AgBr} = 4E−13;
[Ag^{+}]_{3} × [I^{−}]_{3} = K_{AgI} = 3E−16.
For a solution saturated simultaneously with the three silver
salts, the value of the concentration of the silver-ion is the
same in the three solubility-products (see p. 164). Consequently,
for such a solution, with which the three solid salts are in
equilibrium, ‹the ratios of the concentrations of the anions› must
be: [Cl^{−}] : [Br^{−}] : [I^{−}] = K_{AgCl} : K_{AgBr} : K_{AgI} =
3 × 10^5 : 1300 : 1. That is, if silver nitrate is added to a
mixture of iodides, bromides and chlorides, ‹silver iodide› must
be precipitated first, until the concentration of bromide-ion,
in solution, is 1300 times as great as the concentration of the
iodide-ion left in solution. Then ‹bromide› and traces of iodide
of silver will be precipitated, ‹until the concentration of
chloride-ion› is 300,000 times as great as the concentration of
iodide-ion and some 250 times as great as the concentration of the
bromide-ion. In other words, if silver nitrate is added gradually to
such a mixture, iodide-ion and bromide-ion will be almost completely
removed from solution ‹before a precipitate of silver chloride can
be in equilibrium› with the solution. This gives us a convenient and
rapid method of detecting chlorides, if present, [p304] in more
than small quantities, with iodides and bromides. Silver nitrate,
a few drops at a time, is added to the solution and the mixture
vigorously shaken after each addition. As long as a yellow (AgI)
or yellowish (AgBr) silver salt is precipitated on the addition of
silver nitrate to the supernatant liquid (the precipitate settles
quickly), silver nitrate is added as before; when the color becomes
quite pale, the solution is filtered and silver nitrate added a
drop at a time; if a ‹pure white precipitate› results finally,
chloride-ion is present in the mixture (‹exp.›).
«Complex Ions.» Instances of the rôle of complex ions in the
analysis for acid ions are numerous. One of the most interesting
illustrations is the application of the equilibrium conditions
for the complex silver-ammonium-ion (p. 224) to the separation of
silver chloride, bromide and iodide. A rather more convenient and
more sensitive method[597] for detecting the three halide ions in
the presence of each other than the method just considered, may be
discussed from this point of view.
The condition of equilibrium between silver-ion, ammonia and
silver-ammonium-ion is expressed in the relation:
[Ag^{+}] × [NH_{3}]^2 / [Ag(NH_{3})_{2}^{+}] = K = 1 / 10^7.
The concentration of silver-ion, which may exist in an ammoniacal
solution, evidently must decrease rapidly with increasing
concentrations of the free ammonia. Now, let us imagine only
sufficient free ammonia, in solution, added to a mixture of
silver chloride, bromide and iodide, to keep the concentration
of silver-ion, which can exist in the solution, say at
[Ag^{+}] = 6E−9, which is just 1 / 100th of the concentration of
silver-ion in a saturated aqueous solution of silver bromide.
Such a solution of ammonia, in contact with the three silver
salts mentioned, will dissolve silver chloride, if sufficient
is present, until [Cl^{−}] = K_{AgCl} / 6E−9 = 0.017 molar.
At the same time, silver bromide would be dissolved until
[Br^{−}] = K_{AgBr} / 6E−9 = 0.000,06 molar. In other words, ‹silver
chloride could be dissolved in some quantity›, while silver bromide
is dissolved only in traces (the ratio of [Cl^{−}] : [Br^{−}] is
again about 250 : 1). When such an ammoniacal extract is acidified
with nitric acid, almost pure (white) silver chloride would be
precipitated and only traces of bromide would be lost. After the
extraction of the chloride, an ‹increased concentration of ammonia›
would lead, similarly, to a solution in which ‹silver bromide would
dissolve readily› and only traces of the iodide be lost, and thus a
separation of bromide and iodide may be effected.
In Hagar's method, the concentration of ammonia, required to
dissolve silver chloride with but traces of bromide, is attained by
the use of a solution of ammonium sesqui-carbonate,[598] in which
free ammonia is present only in small concentration, as a result
of the hydrolysis of the salt. After the [p305] extraction of
the chloride by this solution, the bromide is extracted with a 5%
solution of ammonia.
«Complex Ions of Acid Ions with Other Acids.»—In the study of
complex ions we found that positive ions (silver, cupric, etc.) may
form complex positive ions with ammonia[599] or complex negative
ions with acid ions (‹e.g.› with cyanide-ion). In the study of the
acid ions we also meet instances of complexes formed by the ‹union
of two acids› to form a new complex acid. Ammonium phosphomolybdate,
an important salt that is extremely useful in detecting the presence
of phosphate-ion, is the most interesting instance of the salt of
such an acid, which is met in elementary qualitative analysis.[600]
Ammonium phosphomolybdate, (NH_{4})_{3}PO_{4}, 12 MoO_{3}, is the
salt of a complex phosphomolybdic acid, formed from phosphoric
acid, O:P(OH)_{3}, and molybdic acid, O_{2}Mo(OH)_{2}, by a loss of
water, much as potassium dichromate is formed from potassium acid
chromate [KO(CrO_{2})OH + HO(CrO_{2})OK ⇄ KO(CrO_{2})O(CrO_{2})OK].
The only difference between the two actions lies in the fact that,
in the case of the dichromate, anhydride formation occurs between
two molecules of a single acid; in the case of the phosphomolybdate,
anhydride formation takes place between molecules of different
acids, and a much larger number of molecules is involved. If
we suppose the combination between the two acids to proceed
symmetrically,[601] we may consider the following to be the action:
O:P(OH)_{3} + 3 [HO(MoO_{2})OH + HO(MoO_{2})OH + HO(MoO_{2})OH +
HO(MoO_{2})OH] ⇄
O:P[O(MoO_{2})O(MoO_{2})O(MoO_{2})O(MoO_{2})OH]_{3} + 12 H_{2}O.
Intermediate complex acids, containing less molybdic acid, are no
doubt formed first (the action is a relatively slow one), and the
action proceeds until the formation of an insoluble salt leads
to the final precipitation of all of the phosphate in this form.
The precipitate shows the characteristic behavior of an acid
anhydride—‹alkalies dissolve it readily› and form phosphate and
molybdate—‹e.g.› ammonium hydroxide forms [NH_{4}]_{2}HPO_{4} and
(NH_{4})_{2}MoO_{4} (‹exp.›). Dichromates, in a similar way, are
converted by alkalies into chromates, an action which may readily be
followed by the change in color (‹exp.›).
«Oxidation and Reduction.»—While fractional precipitation of silver
iodide, bromide and chloride, and fractional solution of the silver
salts in ammonia are convenient methods for detecting the three
halide ions in the presence of one another, the ‹most accurate› and
‹most convenient› methods for this purpose depend on the different
sensitiveness which iodide, bromide and chloride ions exhibit
towards ‹oxidizing› agents. Of the three halogens, iodine shows the
smallest tendency to form its ion (see the table, p. 294), chlorine
the greatest. ‹Vice versa›, of the three halide ions, iodide-ion
is most readily, chloride-ion least readily, oxidized. Treatment
with a mild oxidizing agent, such as ferric-ion [p306] (see Chap.
XIV and Laboratory Manual under iodide-ion), suffices to oxidize
iodide-ion to iodine: 2 Fe^{3+} + 2 I^{−} ⥂ 2 Fe^{2+} + I_{2}.
Bromide-ion and chloride-ion are left practically unaffected
by this agent (see Chap. XIV). A somewhat stronger oxidizing
agent, chromic acid (or its ion Cr^{6+}, see Chap. XV), oxidizes
bromide-ion and leaves chloride-ion practically unaffected:
2 Cr^{6+} + 6 Br^{−} ⥂ 2 Cr^{3+} + 3 Br_{2}. This method of
fractional oxidation forms one of the most convenient and sensitive
methods for detecting the three halide ions in the presence of one
another.[602]
We shall discuss here only one other oxidation-reduction reaction,
taken in connection with the laboratory work—the ‹oxidation of
hydroiodic acid by exposure to the air› and the resistance to
oxidation shown by an ‹iodide›, such as potassium iodide, under
the same conditions. The following method of proximate analysis
of the chief relations involved may also be used to interpret the
contrast in the behavior of hydroiodic acid and that of hydrobromic
or hydrochloric acid (Laboratory Manual, ‹q. v.›). In all of these
cases the actual relations are rendered more complex in consequence
of secondary reactions, than is indicated in the text that follows:
it is intended only to outline the most effective of the factors
involved and to illuminate the qualitative results observed.
«Oxidation of Hydroiodic Acid by Air.»—The oxidation of hydroiodic
acid, or of potassium iodide, by the oxygen of the air may be
considered (Chapters XIV and XV) to involve primarily the action
4 I^{−} + O_{2} + 2 HOH ⇄ 2 I_{2} + 4 HO^{−}. (1)
The condition for equilibrium will be
[I^{−}]^4 × [O_{2}] / ([I_{2}]^2 × [HO^{−}]^4) = K_{equil.} (2)[603]
‹A system in which› I^{−} ‹is directly in equilibrium with› I_{2}
(for which [I^{−}]_{1}^2 : [I_{2}]_{1} = K_{I^{−}, Iodine} = 5.6E29,
at room temperature (p. 298)) ‹and in which, at the same time›,
HO^{−} ‹is directly in equilibrium with› O_{2} (for which at room
temperature [HO^{−}]_{1}^4 : [O_{2}]_{1} = K_{HO^{−}, Oxygen} =
8.2E49 (p. 298)) ‹would also represent a condition of equilibrium
for the› «four» ‹components›. We find thus
K_{equil.} = K_{I^{−}, Iodine}^2 / K_{HO^{−}, Oxygen} =
(5.6E29)^2 / (8.2E49) = 4E9. (3)
With the aid of this constant and of equation (2) we can obtain,
at least, an approximate interpretation of the results of the
exposure of hydroiodic acid and of potassium iodide to the
influence of atmospheric oxygen.[604] We may [p307] calculate,
first, what concentration of free iodine would be ‹required› to
‹prevent› «oxidation» of «hydroiodic acid», in molar solution,
by the oxygen of the air, ‹i.e.› to establish equilibrium.
We will call ‹x› that concentration of I_{2}. As hydroiodic
acid is a very strong acid, ionized to the extent of about 80%
in molar solution, we may, with sufficient accuracy for our
purpose, consider it completely ionized and put [I^{−}] = 1 and
[H^{+}] = 1. Since at 25° [H^{+}] × [HO^{−}] = 1.2E−14 (p. 104),
we may put [HO^{−}] = 10^{−14}. The concentration of oxygen in the
air, at room temperature, may be considered to be approximately
[O_{2}] = (1/5) × (1 / 23.9). Inserting all these given values in
equation (2), we have
[I^{−}]^4 × [O_{2}] / ([I_{2}]^2 × [HO^{−}]^4) =
1 × (1/5) × (1 / 23.9) / (‹x›^2 × (10^{−14})^4) = 4 × 10^9. (4)
Solving for ‹x›, we find ‹x› = 10^{22} = [I_{2}]. That is, free
iodine of this enormous concentration would be required to prevent
oxidation of hydroiodic acid in molar solution by the oxygen of the
air at room temperatures. It is obvious that hydroiodic acid must be
extremely sensitive to oxidation by exposure to air.
One might estimate, in a similar way, the extent to which hydroiodic
acid, of a given concentration, would be oxidized by air before
equilibrium would be reached. The process would involve simultaneous
changes in three factors—iodide-ion is destroyed, iodine is formed
and hydroxide-ion increases, as the result of the neutralization
of hydrogen-ion by the hydroxide-ion formed in the action (see
above). The solution of the equilibrium equation is too involved for
the elementary purposes of this discussion: it leads to the same
qualitative conclusion as was just reached.
«Oxidation of Potassium Iodide by Air.» We may now ask what the
relations would be, if we used a molar solution of potassium
iodide in place of the free acid. [I^{−}] and [O_{2}] would have
the same value as before. The solution being originally neutral,
[HO^{−}] would at first have the value √(1.2E−14) = 1.1E−7. But when
potassium iodide is exposed to the air, if iodine is liberated, the
solution becomes ‹alkaline›[605] (HO^{−} is formed according to
equation (1)) and the concentration of HO^{−} consequently ‹grows
continuously greater›. We will, therefore, formulate the problem as
follows: ‹how much iodine[606] must be liberated, by oxidation of
iodide-ion, in molar potassium iodide solution in order to establish
equilibrium?› For every ‹two› molecules of iodine liberated,
‹four› HO^{−} ions are formed (equation (1)). If we call ‹y› the
concentration of iodine at the point of equilibrium, then 2 ‹y› is
the concentration of [HO^{−}] at that point, formed by the oxidation
process. Inserting the given values[607] in equation (2), we have
[I^{−}]^4 × [O_{2}] / ([I_{2}]^2 × [HO^{−}]^4) =
1 × (1/5) × (1 / 23.9) / (‹y›^2 × (2 ‹y›)^4) = 4E9.
[p308]
Solving for ‹y›, we find ‹y› = 0.007. That is, in molar
solution, about 1.4% of the iodide[608] would be oxidized (carbonic
acid and other acids being excluded); in 5 c.c. (see Lab. Manual,
p. 73) 9 ‹milligrams of iodine[609] would be liberated› to reach a
condition of equilibrium.[610]
It is thus clear that the conditions for equilibrium between a
solution of an iodide and air would be satisfied, in the case of an
alkali iodide, by the liberation of a mere trace of iodine, whereas,
as was previously shown, in the case of hydrogen iodide, a very
large proportion of iodine must be liberated before equilibrium
could obtain. A careful comparison of the two developments shows
that the difference in result[610] is plainly due to the higher
oxidizing power, the higher potential of oxygen (p. 280), in acid
solutions, containing only a minute concentration of hydroxide-ion,
as compared with its efficiency in neutral or slightly alkaline
solution.
FOOTNOTES:
[589] In the few cases when there is interference, it is provided
against.
[590] ‹Cf.› Fresenius, ‹Qualitative Analysis›, p. 520.
[591] The group includes phosphate, borate, fluoride, oxalate,
silicate, arsenite, arseniate, chromate and tartrate ions.
[592] The ions of the amphoteric acids, arsenic and arsenious
acids, and chromate-ion, which is reduced by hydrogen sulphide to
chromium-ion, are found in the systematic analysis for metal ions.
[593] The ions of the amphoteric acids, arsenic and arsenious
acids, and chromate-ion, which is reduced by hydrogen sulphide to
chromium-ion, are found in the systematic analysis for metal ions.
[594] See Fresenius, ‹loc. cit.›, p. 511, footnote, and p. 520.
[595] Other acid ions which would show the group test—precipitation
of a barium salt in a neutral solution—are determined in other
groups, as follows: arsenite, arseniate and chromate ions in the
group of amphoteric acids, etc. (I); carbonate and silicate ions
in the carbonate group (II); and oxalate and tartrate ions in the
group of organic acids (VII).
[596] The constants refer to 18°. The subindices are used to
distinguish the (unequal) concentrations of the silver-ion and of
the halide ions in the different solutions referred to in the text.
[597] Hagar's method. See Fresenius, ‹loc. cit.›, pp. 356 and 378.
[598] See Fresenius, ‹loc. cit.›, pp. 356 and 378, for the
preparation of the solution and for details of the method, and see
Smith, ‹General Inorganic Chemistry›, p. 566, as to the nature of
the sesqui-carbonate.
[599] They also form complex ions with substances related to
ammonia, such as the organic amines.
[600] Ammonium arsenomolybdate is an analogous salt (see Laboratory
Manual, Part III). A similar complex acid, phosphotungstic acid, is
used in alkaloidal analysis.
[601] The exact structure of the complex acid is not known.
[602] In the Laboratory Manual a second, similar method is also
given.
[603] It is considered that water has a constant concentration
in a dilute solution and that for its active components
[H^{+}] × [HO^{−}] is a constant (p. 176).
[604] In the calculation which follows, which is meant merely
for a rough survey, no account is taken of the formation of
complex ions I_{3}^{−}, or of the tendency of hydroiodic
acid to decompose spontaneously into iodine and hydrogen:
2 H^{+} + 2 I^{−} ⇄ H_{2} + I_{2}, a reaction which could also be
studied profitably with the aid of the equilibrium constants for
I_{2} ⇄ 2 I^{−} and for H_{2} ⇄ 2 H^{+}. The value of the iodide
constant is also uncertain (see p. 273).
[605] 4 K^{+} + 4 I^{−} + O_{2} + 2 HOH ⇄ 2 I_{2} + 4 K^{+} +
4 HO^{−}.
[606] The formation of complex ions I_{3}^{−} and other secondary
reactions (formation of hypoiodite, iodate, etc.) are ignored.
[607] ‹y› has so small a value that we may consider [I^{−}]
‹practically› unchanged.
[608] 0.007 I_{2} = 0.014 I^{−}.
[609] A mole of I_{2} = 2 × 127 = 254 grams;
(0.007 × 254 × 5) / 1000 = 0.009 gram.
[610] The tendency of iodine to form hypoiodous acid, iodates,
etc., is not taken into consideration here and involves another
relation.
INDEX
Numbers marked (†) refer to subjects illustrated by experiments,
«heavy» numbers refer to tables.
Acetic acid, 98, 101†, 111–114†
Acid ions, systematic analysis for, 299 ‹et seq.›
groups of, 301
Acids, 81, «82», 100, «104»
Acid salts, 103
Alkali group, 158, 159
Alkaline earth group, 158, 162
Aluminate-ion, 172
Aluminium group, 158, 188 ‹et seq.›, 195
Aluminium hydroxide, 171 ‹et seq.›, 196
Aluminium-ion, 158, 172, 188 ‹et seq.›
Ammonia, complex ions of, 216 ‹et seq.›, 224
ionization of, 87
Ammonium hydroxide, conductivity of, 77
dissociation of, 160
ionization of, 78†, 114
strength as base, 78†, 79†, 168†
Ammonium-ion, 161
Ammonium salts, dissociation of, 34†, 159
effect on ammonium hydroxide, 168†
Amphoteric acids, group of, 302
Amphoteric hydroxides, 171 ‹et seq.›, 187, 188, 196
Antimony, ions of, 174
Argenticyanide-ion, 225
Arrhenius's theory of ionization, ‹see› Ionization
Arsenic acid, 247, 250, 283
Arsenic group, 158, 210, 242 ‹et seq.›
Arsenic, ions of, 158, 174, 242 ‹et seq.›
Arsenic pentasulphide, 247
Arsenious sulphide, colloidal, 125†
Association, 19, 64
Aurocyanide-ion, 232
Avogadro's hypothesis, 33, 34
Avogadro-van 't Hoff Hypothesis, 15, «37»
Azolitmin, «79»
Barium carbonate, as reagent, 193
Barium-ion, 162
Bases, 79†, «81», 105, «106»
Basic salts, 106, 194
Bismuth-ion, 158
Boiling-point, 17, 36, 38
Borates, hydrolysis of, 90†
Bromine, 252†
Cadmium-ion, 158
complex ions of, 224, 228†
sulphide, 209† ‹et seq.›
Calcium-ion, 162
Carbonate group, 302
Carbonic acid, 90, 100
Chemical activity of ions, 72†, 78†, 87†, 232
of acids, 81, «82», 105†
of bases, 79†, «81»
of molecules, 74, 83, 232
of salts, 74†, 107, 116†
Chemical equilibrium, 90 ‹et seq.›, 96†
and direction of action, 112
and path of action, 235
and physical equilibrium, 139 ‹et seq.›
of electrolytes, 98, 111–115†
Chemical equilibrium, law of, 91, 94
and ionization of strong electrolytes, 108
factors of, 95† ‹et seq.›
limitations to, 95
Chemometer, 253
Chloride group, 302
Chloride-ion, oxidation of, 275, 276†
Chlorine, 46, 252†, 275
Chromium-ion, 158, 188 ‹et seq.›
Chromic acid, 287
Cobalt-ion, 158, 192
Cobalticyanide-ion, 229
Cobalt sulphide, 192
Colloidal condition, 125† ‹et seq.›
definition of, 130
relations to analysis, 131, 136†
solution theory of, «128»
suspension theory of, 129
Colloids, electric charges on, 131†
precipitation of, 133†
protective action of, 136, 137†
Complex ions, 88, 216 ‹et seq.›
applications in analysis, 220†, 230†
of acid ions, 305
of ammonia, 216† ‹et seq.›
of cyanide-ion, 225† ‹et seq.›
of halide ions, 237
of oxide-ion, 238
of sulphide-ion, 238, 246
organic, 238†
Concentration, definition of, 91
Conductivity, 44†, 46, 47†, 49, 51, 56
partial, ‹see› Mobilities
Copper, equilibrium with cupric-ion, 258, 264†
oxidation of, 257† ‹et seq.›
solution-tension of, 259†
Copper group, 158, 199 ‹et seq.›, 210, 216 ‹et seq.›
Cupric-ion, complex ions of, 224†
equilibrium with copper, 258, 264†
reduction of, 257†
sulphide, 212
Cuprous-ion, complex ions of, 228†
Cyanide-ion, complex ions of, 88†, 225† ‹et seq.›, 232
Dialysis, 128
Dielectric constants, 62, «63»
Diffusion, of gases, 29†
of ions, 59†
of solutes, 9†
Dissociation of complex ions, 219†
electrolytic, ‹see› Ionization
gaseous, 34†, 36, 96†
Dry salts, 74†
Electrolysis, 46, 58
Electrolytes, ‹see› Ionogens
Electron theory, 42
Equilibrium, ‹see also› Chemical equilibrium and
Physical equilibrium
constants, «95», «98», 233, 236, «298»
oxidation and reduction, 258, 265 ‹et seq.›, 267†, 273 ‹et seq.›
Faraday's law, 58
Ferric hydroxide, 127†, 170†
Ferric-ion, 88†, 251†, 252†, 269†
Ferricyanide-ion, 88†, 230†
Ferrocyanide-ion, 88†, 230†
Ferrous-ion, 88†, 251†, 252†, 269†
Formaldehyde, 289† ‹et seq.›
Fractional oxidation, 305, 306
Fractional precipitation, 163, 165†, 303
Fractional solution, 304
Freezing-point, 17, 36, 38
Fused salts, 75†
Gold, 158, 242 ‹et seq.›
colloidal, 126†
‹see also› Aurocyanide-ion
Groups of metal ions, 157
Groups of acid ions, 301
Heat, 75, 76†
Heterogeneous equilibrium, ‹see› Physical equilibrium
Hydrogen, oxidation of, 277, 278, 280†
Hydrogen chloride, 34, 37, 42, 72†, 73, 84
Hydrogen-ion, action on indicators, 79, 105†
chemical activity, 72–74†, 81, «82», 278 ‹et seq.›
concentration for precipitation by H_{2}S, 213
mobility, 54†, 56
Hydrogen sulphide, ionization of, 199, 245
oxidation of, 251†, 254†
precipitation by, 90†, 199† ‹et seq.›, 203†
Hydrol, 175
Hydrolysis of salts, 127, 178† ‹et seq.›, 190†
Hydroxide-ion, action on indicators, 78†, 79
chemical activity, 79, «81», 168†
mobility, 54†, «56»
Indicators, «79», 165, 214
Instability constants, 219, 224, 226
Iodide-ion, 88†, 272†, 305
Iodine, 273, 305
Ionization and chemical activity, 69, 72†, 79†, 90† ‹et seq.›, 116†,
232
and conductivity, 44†, 77†, 115†
and dielectric constants, 62–«64»
and electron theory, 42
and Faraday's law, 58
and osmotic pressure, 67
and solvents, 61
constants, 98, 100, «104», «106», «108»
degree of, 50
exceptional, of certain salts, 115†
in stages, 100, 102†
of acids, 69, 98, 100, «104», 108
of bases, 69, 105, «106», 108
of colloids, 132
of salts, 69†, 74†, 75†, 107, «108», 115†
theory of Arrhenius, 40, 41, 51
theory of Clausius, 51
Ionogens, 69
Ion-product, ‹see› Solubility-product
Ions, 42
charges on, 41, 58
combination with solvents, 42, 65
composition of, 69†, 89†
migration of, 45†, 53, 54†
mobilities of, «56»
Kinetic theory, of gases, 26
and osmotic pressure, 26
Lead-ion, 158
sulphide, 212, 213
Light, 76
Litmus, «79»
Magnesium hydroxide, 168†
Magnesium-ion, 162
Manganous-ion, 158
Mass action, law of, 93, 96†
Mercuric chloride, 57†, 115†
Mercuric cyanide, 115†, 116†, 231†
Mercuric-ion, 158, 257
Mercurous-ion, 158
Mercury, 257
Methyl orange, «79»
Methyl violet, 213
Mobilities of ions, 53†, «56»
Molecular weights, 33, 36, «37»
Nernst's formula, 261 ‹et seq.›, 296 ‹et seq.›
Nickel-ion, 158
cyanide-ion, 229
sulphide, 192
Nitrate group, 302
Nitric acid, 288†
Organic acids, group of, 302
Organic substances, complex ions of, 238
oxidation of, 289† ‹et seq.›
Osmosis, 10†, 22, 24†
Osmotic pressure, 8† ‹et seq.›
and ionization, 40, 67
definition of, 10
indirect determination of, 16
laws of, «12–20»
measurement of, 10
theories of, 32
Oxalates, complex ions of, 241†
Oxidation, 251† ‹et seq.›, 282† ‹et seq.›
by electric current, 252†
definition of, 251
of organic compounds, 289†
production of current by, 253†
relation to theory of ionization, 251† ‹et seq.›
Oxidation-reduction, reversibility of, 256†, 268† ‹et seq.›
interpretations of, 251, 282, 286
Oxygen, oxidation by, 277, 278, 280†, 305
Oxygen-hydrogen cell, 280†
Permanganic acid, 287†
Perpetuum mobile, 12
Phases, 118
Phenolphthaleïn, «79»
Phosphate group, 302
Phosphomolybdates, 305
Phosphoric acid, 102†, 134
Physical equilibrium, 118† ‹et seq.›
and chemical equilibrium, 139 ‹et seq.›
applications of law of, 120
law of, 118
Platinum, 158, 242 ‹et seq.›
Potassium hydroxide, 77†−81†, «106»
Potassium-ion, 158, 161
Potential differences, 261 ‹et seq.›
and concentrations, 261, 263†
and osmotic pressures, 261
sign of, 261
Precipitates, ‹see also› Precipitation
solution of, 151†
washing of, 148
Precipitation, 122†, 145† ‹et seq.›
and ionization, 74†, 90, 152†, 220 ‹et seq.›
fractional, 163, 165†
influence of a common ion, 144†, «146», «147»
influence of electrolytes, 144†, 150†
of electrolytes, 145†
Primary ionization, 101, 102†
Purple of Cassius, 126†, 134†
Reduction, ‹see also› Oxidation
by electric current, 252†
definition of, 252
interpretations of, 251, 282, 286
production of current by, 253†
Salt-effect, 82, 109, 110†
Secondary ionization, 101, 102†, 246
Silver, colloidal, 127†
Silver-ammonium-ion, 217 ‹et seq.›
Silver bromide, 224, 303
Silver chloride, colloidal, 138†
precipitation of, 150
solubility of, 221
Silver chromate, 141, 165†
Silver group, 157, 199 ‹et seq.›, 216 ‹et seq.›
Silver iodide, 224, 303
Silver-ion, 158
as oxidizing agent, 290†
complex ions of, 216†, 225†
Sodium hydroxide, «81», «106», 173
Sodium-ion, 158, 161
Solubility, 121, 123, «146», «147», 153, «155»
Solubility-product principle, 141 ‹et seq.›
applications of, 145†, 147, 149, 151
derivation of, 139 ‹et seq.›
Solution, theories of, 8, 32
Solution of electrolytes, 151
Solutions, concentrated, 15, 32, 142
dilute, ‹see› Osmotic pressure
non-aqueous, 62, 73†, 84
supersaturated, 121†
Solution-tension, electrolytic, 258 ‹et seq.›
constants, 259, 266, «294», «295»
Solvents and ionization, 61, «64»
and solubility of electrolytes, 154, «155»
Strength of acids, «104»
Strength of bases, 78†, «106»
Strontium-ion, 158, 162
Sulphate group, 302
Sulphide-ion, complex ions of, 238, 246
concentration of, «202»
oxidation of, 251†, 254†
Sulphides, precipitation of, 199†, 203† ‹et seq.›
solubilities of, 203 ‹et seq.›, 212
Sulpho-acids, 244 ‹et seq.›
Sulpho-bases, 244† ‹et seq.›
Sulpho-salts, 243 ‹et seq.›
Sulphuric acid, ionization of, 103†
Supersaturation, 121†
Systematic analysis, of metal ions, 157 ‹et seq.›
of acid ions, 299 ‹et seq.›
Tartaric acid, complex ions of, 238†
Tin, ions of, 158, 174, 242 ‹et seq.›
Unsaturated compounds, 64
Uranyl salts, 286
Valence, 59
and precipitating power of ions, 135
Van 't Hoff's hypothesis, ‹see› Avogadro
Van 't Hoff's theory of solution, 12 ‹et seq.›
apparent exceptions to, 18
Vapor tension, 17, 36, 38
Velocity of action, 92
Water, ‹see also› Hydrolysis
composition of, 175
dielectric constant of, 63
formation by electrolytic oxidation, 280, 282†
ionization by, 37, 41, 47†, 61, 73†, 74†
ionization of, 53, «104», «106», 176 ‹et seq.›
oxonium ions of, 238
secondary ionization of, 246, 278
Zinc, 257†, 266†
Zinc group, 158, 188 ‹et seq.›, 210
Zinc-ion, 266
Zinc sulphide, 204† ‹et seq.›
TRANSCRIBER'S NOTE.
Original printed spelling and grammar is generally retained.
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mathematical and chemical formulas and equations were modified in
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Furthermore, E notation has also been substituted herein for most
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throughout the book, in words such as "hydrogen-ion" versus "hydrogen
ion" or "non-ionized" versus "nonionizied". These have been retained.
The word "difficulty" was sometimes employed as an adverb; herein it
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Page 45: Changed "permangante" to "permanganate".
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in ionization between potassium hydroxide and ammonium hydroxide).".
Page 104: In the table, "Hydrogen sulphide" was changed to
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the sequence of note anchors in the table. The anchor in the table
title originally linked to a footnote instead of a table note; this
footnote was converted to a table note (the first one).
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Constants of Bases" was changed from "3" to "F".
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balanced equations and two unbalanced equations using vertical arrows
has been rearranged into four balanced equations, with horizontal
arrows. Such rearrangements have been silently performed elsewhere.
On page 217, a schema of three equilibrium equations was separated
into three equations, and also removed from its embedded location
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represented herein with a more readily available character "⥂"—the
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Page 125: In the reaction of arsenious oxide with hydrogen sulphide,
changed "H_{3}S" to "H_{2}S".
Page 146: In the table heading, "[CH_{3}COO^{-}]" was changed to
"acetate" to save space. The same thing was done in the table on page
146.
Page 157 "Mendelejeff" changed to "Mendeléeff".
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"glycocoll" a.k.a. glycine was shown roughly as follows:
"H_{3}N.CH_{2}COO" with a line drawn between the "N" and
the right-hand "O". This has been herein modernized to
"H_{3}N^{+}.CH_{2}.COO^{−}", showing explicitly the zwitterion
structure (the second full stop added to emphasize that "COO" is a
unit). The html and mobile versions reproduce the original structure.
Page 197: Changed "[AlO^{3−}] = ‹y›" to "[AlO_{3}^{3−}] = ‹y›".
Page 226 etc.: Both forms "Bodlaender" and "Bodländer" are retained.
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U+2017, Double Low Line, e.g. "‗OH‗".
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esters (p. 81)".
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from "156" to "56".